Induced Reactions and the Higher Oxides of Iron - American Chemical

BI- D.4NFORTH R. HALE

Introductory The induccti reactions in which ferrous iron is the inductor have caused iiiueh pother among investigators ever since 18j8, when Schonbeinl found thai blood corpuscles and ferrous salts accelerate the decolorization of indigo bluc by hydrogen peroxide. In the attempt to account for these reactions, no fever than eleven intermediate compounds and radicals containing iron have been postulated, depending on the oxidizing agent and the acidity of the solution. A few of these have Fimilar empirical formulae, but all are different in thpir assumed structure. The two classifications proposed-the one by Bchiloiv (1903) based on Ppecificity of reagent and the other by Miller (1907) based on effect of reagent on reaction velocity-are unsatisfactory, and seem to have no pragmatic value. Aside from its purely intrinsic interest, the problem is important as bearing on such questions as the passivity of iron, the anomalous break in potentiometric titration curves involving ferrous iron and an oxidizing agent, and oxidation in biological chemistry. The purpose of the research with which this thesis deals is to investigate the more probablc of the postulated intermediate compounds, to sift the evidence for these, and if possible to develop B mechanism for the reaction which will explain the observed phenomena. Perhaps the classic example of an induced reaction is the oxidation of an arsenite solution by nieans of air in the presence of sodium sulphite. Friedrich Mohr csllcd attention to this arsenite reaction in 185j, but to this day its mechanism is obscure. .is this is the induced reaction most often referred to, and one that has the dignity of age, it may well serve here to illustrate the significance of the term. Speaking of the stability of an alkaline tenth-normal arsenite solution in contact with air, lIohr2 said: “When sodium arsenite is allowed to stand in an open vessel with a little sodium sulphite added, it is not long before arsenic acid is also found in the solution. IYhile the sodium sulphite is undergoing oxidation the arsenious acid is infected ivith the oxidation procesy. I placed in a flask some of thp solution that I had already kept ten months without change in strength, and dropped a single crystal of sodium sulphite therein. After two weeks the solution gave a red-brown precipitate with a silver solution, and with magnesium mixture gave in large amounts the characteristic, coarsely-grained, crystalline precipitate of magnesium-ammonium arsenate.” I t will be remembered that sodium sulphite solution is rapidly oxidized to sulphate if allowed to stand in contact, with air. The two reactions may be written separately thus: 0 2 zSa2S03 = zSa2S04 O? zSa3-isOB= zSaa.Is04

+ +

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where the first is the rapid reaction and the second is the reaction which apparently does not occur spontaneously at all, but which occurs readily enough if the other reaction is taking place in the same solution. F. Kesslefl in 1863 developed the terminology which will be used in the sequel: The rapid reaction, the oxidation of the sulphite is the p r i m a r y reaction; the slow or nil reaction, the oxidation of the arsenite, is the secondary TPaCtiO?l, which is said to be induced by the occurrence of the former in the same solution. The oxygen, logically enough, is the actor; the arsenite is the ncceptor; and the sulphite, the actual accelerating agent, is the inductor. The proportion in which the actor divides itself between the acceptor and the inductor is called the induction factor, the relative amounts being measured not in grain molecules but in grain equivalents. The story of the reactions induced by ferrous iron is hut a chapter in the history of the induced reaction, which, first recognized about 18j5,has been the subject of much investigation and of various theories and attempts at classification, resulting, however, in so small a body of definite knowledge and proven theory that the term is rarely found to-day in the indices of chemistry textbook?. Induced reactions have been likewise designated as “coupled” (W.Ostwald)6 and “sympathetic” ( J . W. Mellor)’ reactions; but these terms are quite as rarely met with. Wagners calls the phenomenon “pseudocatalysis.“ Ilhar’s translationg of Ostwalcl‘s term is “catalysis by transvection. ’‘ C. 1.; Schonbein’O was probably the first experimenter to speculate on the actual mechanism of induced reactions. In 1 8 j 8 he showed that ozone is formed when phosphorus is oxidized by oxygen, and that when benzaldehyde or turpentine is oxidized, indigo in the same solution is simultaneously oxidized to colorless isatin. He demonstrated that the oxygen which is eonsumed is divided equally between the two substances oxidized, and this has been verified by van‘t HofP1, Jorissen’*, EnglerL3, and others for several organic and inorganic substances. Schonbein’s explanation” ( I 860) mas that during the oxidation an oxygen molecule breaks up into two atoms of opposite charges, and that these naturally would tend to oxidize different substances.’:, lfi Engler’3 calls this phenomenon “autoxidation” or “auto-0sidation”actually a misnomer, since, although the reaction is an oxidation, neither of the substances oxidized has the power of oxidizing itself as, for instance, potassium hydrochlorite has. The terms, however, are in the literature and the dictionaries, and must be reckoned with. At any rate the phenomenon is an induced reaction. In the phosphorus experiment the oxidation of the phosphorus is raDid; the oxidation of gaseous oxygen to ozone is relatively extremely slow or nil, but is greatly accelerated while the phosphorus oxidation is in progress. hloritz Traube” in 1882 expressed disagreement with the positive and negative oxygen atom theory. He showed that many oxidations, if occurring in the presence of mater, induce the formation of hydrogen peroxide in small quantities. An instance of this is the reaction of zinc and oxygen,

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in which, by Kessler’s nomenclature, the actual oxidation of the zinc may be called the primary reaction. Then zinc is the inductor and water is the acceptor of the induced reaction. Traube suggests that molecular oxygen adds itself to the hydrogen of a water molecule, forcing out the oxygen atom, which then combines with the zinc. If this is the correct explanation, then the oxygen first unites with the acceptor. In 1897 A . Bach18 and in 1902 Wilhelm h l a n ~ h o t ‘stated ~ the view that the substance undergoing rapid oxidation, the inductor now, unites with oxygen to form a peroxide. It is this theory, particularly for those reactions in which ferrous iron serves as inductor, which has received most attention and which is supported, for certain cases, in the present thesis. Manchot, of all the many investigators in this field, has probably done the soundest, practical and theoretical, work on induced reactions. The reactions he studied and the compounds he post,ulated have served as the basis for a number of researches, including this one. His first paperZ0appeared in 1899; his latest?’ in 1 9 2 j . It. Luther and N.Schilow2*in 190.3 developed a general classification of induced reactions which was based on the specific action of inductor, actor, acceptor, and combinations of any two. Thus the induced reaction involving ferrous iron comes under Class -1-inductor specific-for if suitable other compounds are used for the actor and acceptor the reaction is still induced, but if another metal is used in place of iron the reaction will probably not go at all. This system gave six classes, but it is very one-sided, for all reactions involving ferrous iron as the inductor come under Class A, and the eight known reactions in which sulphur dioxide is inductor are distributed among the other five classes. This is not necessarily an objection to the classification, because the reactions might perhaps be distributed that way. The real objection is that Luther’s classification is quite arbitrary. W. Lash Millerz3of Toronto wrote in 1907 a paper which was the culmination of several investigations carried on under his direction on the subject of reaction velocity in oxidation-reduction processes. These are published in the Journal of Physical Chemistry beginning in 1903. He himself puts forward no theory; but, in an early paper by Miss Clara B e n ~ o n a, ~com~ licated theory for the iron reaction was advanced which involves the assumption of a “ferroiodion,” FeIf, a device which serves to account for certain peculiar effects occurring in the reaction velocities. The essence of her findings is that in the absence of iodides the rate of oxidation of ferrous iron by chromic acid is proportional to the square of the acid concentration and to the square of the ferrous salt concentration, whereas in the presence of iodides this rate of oxidation is proportional to the third or fourth power of the acid concentration, and to the first power of the ferrous salt concentration. This result seems to throw- doubt on hlanchot’s primary reaction, the oxidation of ferrous iron to a peroxide, since this primary reaction ought to take place in the same manner whether iodide is present or not. By assuming the temporary existence of the ferroiodion, the change in velocity when

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iodide is added becomes plausible, but Miss Benson does not attempt to explain the particular exponents that she found. These reaction velocity measurements will be mentioned again later. Miller classified induced reactions using as a basis the effect of the reactants on the reaction velocities. I n Class I the presence and concentration of acceptor has no effect on the rate at which the inductor is acted upon. In Class I1 the rate of change of the actor is independent of the presence and concentration of the acceptor; and in Class 111 the addition of acceptor changes the effect of the reagents on the rate a t which the inductor is destroyed. Some time in 1923, Wilder I). Bancroft?j outlined a classification of induced reactions based on the possible mechanisms through which they might occur. This seemed more logical than a classification founded on arbitrary properties of the reactions or on reaction velocity. To quote his words: “Let A be an oxidizing agent (actor), which will not react with a reducing agent (1 (acceptor), but which reacts with a reducing agent B (inductor) and with C in presence of B.” Five types of induced reactions may then be distinguished: I. B may catalyze the reaction between A and c‘. C may react with a low stage of A. 2. 3. C may react with a high stage of B. 4. B and C may form a complex. j . Combinations of the above may occur. I t should be noted that two varieties of catalysis are admitted to the class of induced reactions by this classification. The inductor is not usually a catalyst, for it disobeys one of the criteria: no permanent chemical transformation. The first type of induced reactions, however, are those in which the inductor acts in a truly catalytic manner in bringing about a reaction between actor and acceptor, but in which a molecule of inductor, having catalyzcd the reaction of any actor and acceptor in its sphere of influence, is transformed by the actor into a substance not possessing the catalytic activity. A second form of catalysis belongs under Type three. As an example of this, furfural is not oxidized in the presence of sodium chlorate alone, but is rapidly oxidized if a little vanadium pentoxide is added.$ Vanadium pentoside alone oxidizes furfural and shows the following changes of color: V20j (yellow), Y2O4 (blue), \-?03(green), which are immediately reversed when a crystal of chlorate is added. Assuming the use of an excess of chlorate, the pentoxide is a catalyst and it obeys the criterion of not undergoing permanent change. Yet the reaction is an induced reaction, for the slow furfuralchlorate reaction is accelerated in the presence of the rapid furfural-pentoxide reaction. I n the chapter on Physical Chemistry in the Golden Jubilee Kuniber ( 1 9 2 6 ) of the Journal of the American Chemical Society, Rancroft says: “At the Toronto Laboratory, under the direction of Lash Xiller, work was done for several years on reaction velocity in systems containing oxidizing

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and reducing agents and on systems involving coupled reactions. Miller showed the inadequacy of Luther’s classification of coupled reactions, without being able, however, to substitute a clear and workable classification of his own. This particular problem is one which will be solved long before the Centennial celebration.” At Baker Laboratory a start toward the study of induced reactions was made in I 9 2 7 when IT,G. 1-annoy submitted the reaction, chromic acidtartaric acid induced by arsenious acid, to an investigation and found that the limiting induction factor (correction being made for extraneous influences) had not been carefully determined by previous workers, and found also that the reaction is most probably a member of Type 4 of Bancroft’P classification, a complex being formed between inductor and acceptor.26 .4s a phenomenon related to the induced reaction, it may be mentioned that the passivity of iron in the presence of strong oxidizing agents has long been thought to be caused by a thin film of adsorbed higher oxide. In 1;90 ,James Keirz7wrote: ‘’I put some pieces of clean fresh iron wire into a concentrated and red fuming nitrous acid. No apparent action ensued; but the iron W R S found to be altered . . . that is, it was rendered incapable of being attacked either by a phlogisticated solution of silver or by dephlogisticated The alteration thus produced on the iron is very superficial. The least rubbing exposes some of the fresh iron beneath the surface, and thus subjects it to the action of the acid.” Faradayz8 in 1836 expressed more definitely the view that the passivity is caused by a film of oxide. Bennett and BurnhamzQin 1917 concluded that adsorbed higher oxide films were formed on passivated chromium, iron, and lead. “Passivity in all cases is the coating of the metal, by adsorption, with a film of a higher oxide which, being more noble than the metal, protects it from the action of the solution.” They were able to show that iron dipped in a perferrate solution (iron ignited with potassium nitrate) became passive. U. R. Evans30 has quite lately isolated the film and calls it ferric oxide, but he has no justification for this. The film producing passivity must be a film stabilized by adsorption, as Bennett and Burnham showed, and such a film can hardly exist by itself when the actual surface upon which the adsorption occurs is removed. Evans’ proof is that the film yields a blue color with ferrocyanide, but if his film were a higher oxide it would give the test for ferric iron, since ferrocyanide would reduce it to the ferric state. As a matter of fact the film will change t o ferric oxide as soon as the substrate i. removed. Freundlich, Patscheke, and Zocher”’ have succeeded in precipitating an exceedingly thin iron mirror on glass, and they observe that the optical properties of this film change on contact with air. I t may be possible to show that this change in optical properties does not correspond to the change in optical properties when iron is oxidized to the ferric state, and thus to prove that a higher oxide is formed, E. S. Hedges4 has discovered a property of passive iron which is also a property of ferric oxide, indicating that the two may be identical. When passive iron is heated in solutions of nitric acid ranging in concentration

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from 90% to 100% (density 1 . 4 2 ) the first yellowing of the liquid due to the solution of the iron occurs a t a temperature of 74.5’ to 75.0”; and ignited ferric oxide starts to dissolve in 100% nitric acid a t 7 jo,and in 90% acid at 72O-77’. However, there are so many properties of the passive iron film not shown by ferric oxide, that these experiments can not offer a proof of the identity of these two substances, but merely indicate the vagueness which is found throughout our meager knowledge of this phenomenon. I n the electrometric titration of dichromate with ferrous iron, the oxidizing potential unexpectedly first rises, and then drops suddenly when nearly the equivalent amount of ferrous salt has been added. This rise of potential becomes intelligible on assuming the intermediate formation of a higher oxide of iron having a potential above that of the dichromate. G. S. Forbes and E. P. BartletP in rg13 first noticed the phenomenon. The observed increment of potential was 0 . 2 volts. Permanganate did not produce this effect, and chlorides were found to hinder it. N. H. F ~ r m a nin~ ~his study of “Bimetallic Electrode Systems for Potentiometric Titrations,” found a pronounced peak in the curve for the titration of bichromate in sulphuric acid by the potentiometric method, using a Pt-Au system. Furman also obtained sharp peaks when permanganate is the oxidizing agent. It will be mentioned later how this phenomenon has been applied to the study of induced reactions by Goard and Rideal. Baudisch and W e l have ~ ~ called ~ attention to the importance in biological chemistry of induced reactions in which ferrous iron is the inductor. The idea that iron salts act as catalyzers in biochemical processes, and play an important part in respiration, particularly, has gained in significance lately through the work of various investigators. The mechanism of this action of iron remained obscure, however, and one was content with the simple assumption that the variable oxidation qtages were responsible for the specific catalytic power of the iron salts. “This general assumption seems, however, t o be of very doubtful validity when one considers that on a change from the ferrous to the ferric state, a very stable form is assumed which can be brought back to the ferrous state only with the use of very considerable amounts of energy. The reduction of trivalent to divalent iron is very difficult to bring about in the ordinary ionic salts of iron. Iron bound in a complex ion, on the other hand, is in general easily brought into each of the oxidation stages. . . . “It is known with considerable certainty that the catalytic action of the iron in respiration and many other biologically important processes is especially related to the properties of the ferrous atom or ferrous ion and not so much to the ferric atom or ferric ion.” D h a F points out that when iron is administered aF a drug it very probably acts though induced reactions: “Iron has long been used in the treatment of anaemia more especially of the form known as chlorosis and it was assumed tacitly that it was readily adsorbed from the alimentary tract and was utiliied by the tissues to form haemoglobin. . . . If the iron is adminstered in the ferrous state, it passes into the ferric condition in the body and

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usually exists as a part of a complex radical and in a colloidal condition. This complex, by coming into contact with the peroxide formed from the inhaled oxygen, forms a higher oxide of iron which oxidizes food materials.”

Hydrogen Peroxide as Actor Historical Schonbein, as has been mentioned, early observed the inductive character of the reaction1 involving hydrogen peroxide, ferrous iron and indigo blue. He proposed the reaction as an analytical test for hydrogen peroxide, and he studied the effects of other oxidizing agents, but his work on iron was mostly of a descriptive character only. I n 1894 H. J. H. Fento@ used ferrous iron to induce the reaction between hydrogen peroxide and tartaric acid. Being an organic chemist he devoted his time to studying the properties of the product, dihydroxymaleic acid. He did not try to work out the quantitative relation in which the iron stood to the amount of dihydroxymaleic acid, nor did he waste many words on his simple theory. He wrote of his discovery as follows: “When tartaric acid in aqueous solution interacts with certain oxidizing agents in presence of a trace of ferrous salt, a solution is obtained which gives a beautiful violet colour on the addition of caustic alkali. . . . Ferric salts are quite inoperative in bringing about the change; but if, in the first instance, the quantity of ferrous salt is very small, the colour produced by the alkali is greatly intensified by adding a few drops of ferric chloride.” This is obviously an induced reaction since ferrous but not ferric salts produce the effect. I n 1898 he wrote.n “Whea tartaric acid is ouidized in presence of a small quantity of ferrous iron, one molecule of the acid loses two atoms of hydrogen, giving rise to dihydroxymaleic acid. The most effective oxidizing agent for the purpose is hydrogen peroxide, but the result is also brought about by chlorine, hypochlorites, bromine, etc., and by atmospheric oxygen in the presence of sunlight. The presence of ferrous iron is essential, but its proportion seems to bear but little relation to the yield of acid in the ordinary course of preparation, the action being, in fact, what is usually termed catalytic. It is necessary that the addition of the iron shall precede that of the oxidizing agent.’, Then in another paragraph : “Dihydroxymaleic acid readily reduced ferric salts in the cold, experiment indicating that two atoms of iron are reduced by one molecule of the acid, so that the ferrous iron is regenerated at the expense of a portion of the acid.” Thus the induced reaction is unusually complicated, and possesses the characteristics of a catalytic reaction when it once gets started. It is probably a series of three consecutive reactions involving two intermediate products: the higher oxide of iron, and the dihydroxymaleic acid. The first intermediate reacts with tartaric acid to give the second, and part of this reduces the ferric iron that has just been formed. It is then more or less of an accident that sufficient of the dihydroxymaleic acid is let undecomposed tomake thisavaluable method of preparingthe latter.

a

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Fenton applied the reaction to the oxidation of many organic acids and polyhydric alcohols, finding the mixture of ferrous sulphste and hydrogen effective in some cases and ineffective in others.3s According to his hypothesis, the ferrous iron displaces two hydrogen atoms and forms a link between the two carbon atoms attached to the central hydroxyl groups; then, under the action of the oxidizing agent, the iron becomes trivalent, and now may be assumed to break away from the organic molecule, being no longer suitable for the link; and this breaking off has released two unsatisfied carbon bonds, which allows the oxygen to enter. Sinec the reaction is catalytic, this explanation or that involving an intermediate higher oxide of iron ~ o u l dfit equally well. If a method could be found for preventing the reducing action of the dihydroxymaleic acid, so that the reaction would be truly inductive, then the ratio of oxidized iron to oxidized tartaric acid could be obtained. This ratio would be one atom to one molecule if Fenton’s explanation is correct, or one atom to three molecules if thr intermediate is FeOa, and so on. Walton and Christensen,ZQwriting on “The Catalytic Influence of Ferric Ions on the Oxidation of Ethanol by Hydrogen Peroxide,” persistently make the error of calling a mixture of ferric salt and hydrogen peroxide, “Fenton‘s reagent.” Fenton himself notes that if the ferrous iron be replaced by ferric in the experiments he performed, or if ferrous iron and hydrogen be mixed in the absence of the organic liquid and be then added to the latter, the effect is markedly different. “Fenton’s reagent,” then, if this mixture is to have a special name, is a mixture of ferrous salt solution and hydrogen peroxide, which however is not mixed until the substance to be oxidized is added to one or the other of the constituents of the reagent. Wilhelm ?vlanchot40at the Chemical Institute of Gottingen investigated in 1901 the reaction velocity of the oxidation of ferrous iron by oxygen, with the purpose of ascertaining to which order the reaction belonged. Knowing that ferrous ammonium sulphate oxidized very slowly in air, and that the ferrous salts of many organic acids react rather rapidly, he put measured amounts of ferrous solution and potassium- oxalate or citrate or tartrate solution into separate compartments of his apparatus, and ran one into the other and started the shaking motor a t the same time that he pressed the stopwatch. His apparatus was filled with oxygen and connected to a burette, the level of whose confining liquid indicated at any time just how much oxygen had been absorbed. He obtained neither a first nor a second order constant, and he concluded that he was observing a complicated reaction. He noticed, however, that in these experiments there occurred a greater absorption of oxygen than corresponded to the ferrous content. Suspecting that this might be due to an activation of the oxygen by the ferrous iron, with concomitant oxidation of part of the organic acid, he conceived the idea of oxidizing the iron in the presence of a large amount of arsenious acid, in which case the activated oxygen ought t o go quantitatively to the oxidation of the latter. He found

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that the mixture absorbed practically double the amount of oxygen necessary t o oxidize the iron to the ferric state, and from this fact deduced the existence of an instable iron dioxide, !?eo2. Continuing his researches in oxidation, Manchot tried the effect of different oxidizing agents on ferrous iron in the presence of an acceptor, and found that when he used chromic acid, hydrogen peroxide, or permanganate, the results could he explained by assuming an unstable higher oxide of formula FenOi. Among other oxidizing agents he tried hypochlorous acid, and here he met with some difficulty, since it oxidized the acceptor rather readily in the absence of an inductor. Tartaric acid proved to be the best acceptor of several examined; but with this he found difficulties in analyzing the reaction mixture. In t,he end, he found that his data pointed to FeOo as the intermediate compound. A different method was used each time for calculating the three formulae, I'eOn, Fe20j, Fe03. The formula FenOa(hydrogen peroxide, chromic acid, or permanganate as actor) was obtained as follows: The formation and subsequent deconiposition of the intermediate higher oxide can be expressed thus: 2FeO ( n + I ) O = FezOsin = Fe?03 no.

+

+

To a large excess of acceptor and actor a small known quantity of ferroui iron was added, such that all the iron would be oxidized quickly to the ferric state. The amount of decomposed acceptor was determined, calculated in terms of equivalents, and then the number of equivalents of oxygen activated by one equivalent of ferrous iron was calculated, which is the n of the above equation, The formula FeO2 (oxygen as actor) was obtained more directly: The actual amount of oxygen absorbed was known, and this went ultimately t o ferric iron and to acceptor. The ferrous iron was used in such small quantity that it was completely oxidized in a short time. Thus the amount of oxygen going to the acceptor for every equivalent, of iron could be calculated in equivalents, the results being the n of the above equation. The formula Fe03 (hypochlorous acid as actor) could not be obtained ,so simply because of the difficulties previously mentioned. By a method of trial and error, Nanchot found that when ferrous salt and hypochlorous acid in the ratio of I to 4 are brought together in the presence of a large excess of tartaric acid, the actor and acceptor are practically completely consumed. Here the four equivalents of actor are the n + r of the previous equation. Other oxidizing agents examined were so slow in their action as to make it impossible to determine what proportions of acceptor and inductor were simultaneously oxidized. These were persulphuric acid, chloric acid, bromic acid, iodic acid, and nitric acid. Although considerable effort was made to isolate these three hypothetical compounds, all early attempts proved futile. There are, however, some data on compounds containing iron with a valence higher than three. I n the compounds of formula M9Fe03in tcrminology used by J. Sewton Frie~id,~'

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which the iron appears to be quadrivalent are called ferrates, and compounds of formula MeFeO4 in which iron appears to be hexavalent are called perferrates. The existence of the perferrates was established about 1841 by E. FremyJ4> J. Denham and Heinrich Rose,44and much of this work was repeated and confirmed 46, 47* 48. Ferrates were prepared in 1909 by L. Moeser and H. Borck,(* and by G. Pellini and D. MeneghinisSo Moeser and Borck prepared the barium and strontium salts using a temperature of 600' and a current of oxygen as the main conditions of experiment. Their analytical methods gave then the empirical formulae, which they wrote as FeOz.nSrO, FeOZ.BaO, etc. Pellini and Meneghini obtained the actual peroxide-not a salt of the acid-by the action of hydrogen peroxide on ferrous and ferric compounds. They employed the special conditions of an alcohol medium and a temperature of - joo to - 70°, and they obtained a reddish precipitate which, analyzed for iron and oxygen, gave a ratio averaging I : I . ~indicating , a formula of FeO? or FezOa. However they regard the FeOz as being analogous to barium peroxide, and as containing bivalent iron; and the iron in the Fe204is COILsidered to be trivalent. D. K. G0ralevich5~ (1926) fused ferric oxide, potassium hydroxide, and potassium nitrate together and obtained K2Fe04,or, with doJble the quantity of nitrate, K2FeOj. H. E. Williamssz (1915)gives several references to quadrivalent iron in the compound KzFe(CN)6.

Experimental Using chromic acid as actor and potassium iodide as acceptor, Manchot's experiments were repeated in this Laboratory, and the results obtained checked his own within moderate error. The important value is the amount of acceptor oxidized, which, the product being free iodine, is readily determined by titration with sodium thiosulphate solution. This value divided by the quantity of ferrous iron oxidized is the induction factor. I n Manchot's experiment^,'^ the factors were 1.83,1.84,I . 8 j , 1.90,and in the repetition they were 1.62, 1.7, 1.83, 2 . 0 . Definite amounts of reagents-potassium bichromate, potassium iodide, and sulphuric acid-were added to z joo cc. of cold distilled water, and the system was brought to oo while being vigorously stirred. A given volume of ferrous solution was added during a minute and a half, then starch was added, and the liberated iodine was a t once titrated with standard decinormal thiosulphate. The two sets of data were not obtained under identical conditions, variations undoubtedly occurring in the mean temperature, in the efficiency of stirring, and in the acidity (the acid was in each case only approximately two-normal). Experiments on the accuracy of the iodine estimation in these solutions indicated that the error was certainly less than four percent, and probably not greater than two.

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-4few reaction velocity experiments showed that within five minutes of the time the ferrous sulphate was added, the reaction was 97% completed. In half an hour equilibrium was practically reached, although a slow liberation of iodine continued indefinitely, indicating the occurrence of a catalytic reaction. This circumstance increased the difficulty of determining the exact end-point of the induced reaction, and thus increased the probable error. From the data it appears that the limiting value of the induction factor is 2 , which corresponds to an intermediate compound of formula Fen05. Now Schilow had shown that the induction factor varies with the acceptorinductor ratio, and even Manchot, using a very few values of this ratio, found that the induction factor varied, although apparently not in a regular manner. In the present investigation the amounts of actor and acceptor have been kept constant, and the amount of the ferrous sulphate-the inductor-has been varied within wide limits; thus the acceptor-inductor ratio has been varied widely. It was found that the induction factor is a function of the variable acceptor-inductor ratio. Hydrogen peroxide was used as the actor instead of bichromate, for it acts as an oxidizing agent whether acid is present or not. I n an acid solution, as Manchot himself noticed, the induced reaction is retarded and the catalytic action of the ferric iron is increased. A few experiments with the reaction velocity in solutions of approximately the same acidity as used in the bichromate reaction showed that the induced reaction was virtually complete in two minutes, perhaps less, but that the iodine liberation continued so rapidly that in thirty minutes the volume of thiosulphate equivalent to all the iodine became twice the volume equivalent to the iodine for the induced reaction itself-in short, that a relatively rapid Catalytic reaction was in progress. This iodine liberation was, indeed, considerably more rapid than that observed in the experiments with bichromate of approximately the same concentration as the hydrogen peroxide, measured in equivalents. With regard to the inhibiting effect of acid on the induced reaction itself, Goard and RideaP suggest that “this retardation of the iodine separation may be accounted for by supposing the initial formation of FezOs to be retarded by the presence of hydrogen ions, since the ferrous ion is well known to be stabilized by acid.” In striking contrast to the behavior in the presence of acid, in a neutral solution or in a solutioh which is just sufficiently acid to prevent the hydrolysis of the ferrous ammonium sulphate, the reaction is remarkably rapid and clear-cut. The solution, after the iodine has been removed with thiosulphate, is a light yellow, precipitating a fine yellow oxide or hydrated oxide of iron after a few j6 Also, after a few hours, the solution turns a very faint blue, which is decolorized by a fraction of a drop of decinorrnal thiosulphate. This slight reaction is probably a direct interaction of hydrogen peroxide and iodide catalyzed by the ferric iron. In two experiments involving typical proportions of reagents, 0.29 cc. of decinormal thiosulphate

1644

DASFORTH R. HALE

were equivalent to the iodine liberated after the induced reaction in each experiment during ten or eleven hours. The total titres at the end of this time were 2 . 5 2 and 6.34 cc. thiosulphate. Thus in neutral solution* the induced reaction yields an unambiguous end-point which permits a more careful examination of the reaction, This circumstance, and the observed variation of the induction factor with thc acceptor-inductor ratio, suggested the following Feries of experiments. Approximately tenth-normal solutions of hydrogen peroxide, ferrous arnmonium sulphate, and sodiuni thiosulphate, an approximately two-tenth normal solution of potassium iodide, and starch ( I g of soluble starch to I O O cc. water) were employed. Neasured amounts of iodide and peroxide solutions were run into one and a half liters of distilled water in a two-liter Pyrex beaker. About 15 cc. of the starch solution mere added, the volume mas made up to two liters JTith distilled miter, and the whole was stirred vigorously for a minute or two. The system was kept at a room temperature of 22' to 24'. The ferrous sulphate solution was added from a burette, usually at a rate such that the drops fell almost too fast to be counted. 1-arying this rate from drops that fell slowly and could be counted with ease, to a stream showing almost no division into drops, made scarcely a measurable difference in the amount of acceptor decomposed. The stirring device was a Witt stirrer with a bulb about 4 cin. long, driven by an electric motor having a disk friction drive for regulating the speed. The bulb of the stirrer was placed close to the bottom and near the side of the beaker, and Lhe speed was regulated so as to be as great as possible without causing air to be drawn down inio the solution by the vortex. This speed was, however, much increased during two of the blank runs. As soon as the ferrous solution had been added the iodine was titrated with thiosulfate. Experiments showed that the induced reaction was complete almost as soon as the last drop had been stirred in. Through the help of a sheet of white paper placed behind the beaker, the iodine titre was obtained accurately to two drops of tenth-normal thiosulphate. Blanks were run as follows: The direct reaction between iodide and peroxide was examined by diluting to two liters 1 2 cc. of the former and 2 cc. of the latter, and stirring. I n forty minutes the liberated iodine was equivalent to 0 . 2 j cc. of S,!ro thiosulphate, and in 2 1 9 minutes it was equivalent to 0.98 cc. I t will be rcmenibered that the induced reaction is complete in not over one minute. The possible oxidation of the iodide by the oxygen of the air was tested by diluting to two liters 1 2 cc. of iodide, and stirring. KOcolor was observed at the end of one and a half hours in either of two experiments, the reaction mixture of one being stirred in the normal manner, and that of the other being stirred so vigorously as to keep air bubbles moving throughout the Tvhole volume of the solution. * "Neutral" is here taken to mean "no acid added"; actually, due to the hydrolysis of the ferrous salt, the solution shows a very slight acidity.

INDUCED REACTIOSY ASD THE HIGHER OXIDES O F I R O S

164.i

That ferrous iron does not catalyze the oxidation of iodide by atmospheric oxygen was s h e - n by diluting to two liters 1 2 cc. iodide, I cc. ferrous solution, and several times the equivalent quantity of sodium fluoride, the purpose of the latter being to keep ferric iron out of the mixture through the formation of the ferri-fluoride complex compound. S o color was obserl-ed at the end of 165 minutes, whether the solution was normally stirred or vigorously aerated. In connection with the great difference between the behavior of the reaction in acid and in neutral solution it was mentioned (vide ante) that in the, latter medium ferric iron does not catalyze the reaction between peroxide and iodide, or between oxygen and iodide, except to a negligible extent.

TABLE I BI used, 1 2 cc. S,'I H 2 0 2used, 2 cc. X 1 I Total volume, 2000 cc. Espt.

Temp.

8

22.4

2

3

Fe" used

Iodine

Fe,' consumed

0.1

0.286

0.1

23

0.2

0

2.i3

60 40

0.72

0.3 0.3

2.13

9

0.3 0.3

547 0.638

0 . 2

24

2.4

40

4

0.5

0.991

0 . j

I .98

24

5

0.i5

1.19

0.;j

1.59

16

'

Induction Factor

2.86

=

Fe" I20

6 16

I

.o

1.19

0.81

1.47

I2

I

.o

0.965

I

.o

0.96j

I2

15

I

.6

1.012

0.99

I .02

7.5

0.56 0.97

6 4 3.4 3

23.8

2.0

0.72

I .28

22

I3

3 .o 3.5

0.961 0.955

I .04

I4

4.0

0.721

I

7 I?

I .OI

.zs

0.92

0.562

To simplify the statement of the concentrations, the aniounts of iodide, peroxide, and ferrous iron are given as the product' of the added volume and the normality; thus these quantities represent the volumes of one-normal solutions diluted to two liters, the total volume of solution in each case. I t was desired to keep the concentration of peroxide and ferrous iron in the neighborhood of one-thousandth normal, corresponding to z cc. of K / I solution in the two liters. Preliminary experiments showed that more potassium iodide than 1 2 cc. S , I solution did not increase the induction factor when I cc. S ' I ferrous solution and 2 cc. K / I hydrogen peroxide were used. These quantities of iodide and peroxide were kept constant in this series of experiments, and the volume of K,'I ferroos solution was varied from 0.1to 1 cc. In Table I the nunibers under "Iodine" represent the volume in cc. of normal iodine solution; they are obtained by multiplying together the volume

1646

DANFORTH R. HALE

of thiosulphate required in the titration and its normality. The ferrous iron is regarded as being combletely oxidized if there is present an excess of hydrogen peroxide after the induced reaction is over. Thus in the first six experiments the quantity of ferrous iron oxidized is the quantity of ferrous iron added. When the combined volume of normal ferrous solution and normal iodine solution exceeds 2 , the volume of normal hydrogen peroxide employed, the ferrous iron oxidized is found by subtracting the quantity of normal iodine solution from 2 , since in this case some ferrous iron remains unoxidized. Although the fluctuations in the induction factor as the ferrous iron concentration increases leave much to be desired, the graph of the data makes unmistakable a conclusion which may be stated in tnis form: allowing for all the factors by which the quantity of iodine liberated depends on the concentration of ferrous iron, the lattpr, when added to an excess of potassium iodide and hydrogen peroxide at great dilution, conditions the decomposition of three equivalents of acceptor per mol of iron oxidized from the ferrous to the ferric state. Applying Manchot's method of determining t,he formula of the intermediate compound, this result points to the formula FeOa, though the data are as adequately explained by the assumption of an intermediate of formula Fe0.zH202. The falling off of the factor as the ratio of acceptor to inductor decreases, or as the amount of ferrous iron increases, is most probably due to the interaction between the intermediate compound and unused ferrous salt: FeO intermediate -+ FesOs. This reaction would obviously increase in velocity with increase in ferrous iron concentration, and to the extent that this occurs, the ferrous iron transfers to t,he acceptor a smaller amount of oxygen than the theoretical maximurn. This side reaction becomes less and less important as the ratio of acceptor to inductor is increased, as is confirmed by the data: a ratio of 30 to I , for instance, corresponds to one of 2.8. Manchot obtained a factor less t'han 2 because he used too much iron-or too small a ratio of acceptor to inductor.56 He himself admits that the side reaction occurs, and on this basis he explains the fact that his induction factors are not integers. I n his experiments on this induced reaction he uses I O cc. of decinormsl ferrous ammonium sulphate, and 60 cc. of N/s pot.assium iodide. This corresponds t o an acceptorinductor ratio of 1 2 to I , and the induction factor he obtained was 1.8. In the present investigation a ratio of 1 2 to I gave an induction factor of approximately 1 . 3 . The reaction between intermediate and ferrous iron has been mentioned by others, and receives considerable support from the fact that iron cannot be passivated in the presence of a ferrous salt. It has been demonstrated to occur by Manchot,21 who showed that when dilute hydrogen peroxide is added dropwise to dilute ferrous sulphate, a situation that would greatly favor this reaction, an end-point is reached when hydrogen peroxide and ferrous sulphate are consumed in the proportions to form a ferric salt. He determined the end-point of the reaction potentiometrically. Using the same solutions and the samc method for obtaining an end-point, he showed ?"hat when the

+

INDUCED REACTIOSS AND THE HIGHER OXIDES O F IRON

1647

ferrous salt is added to the peroxide, this order hindering tho reaction between intermediate and ferrous iron, the end-point’ occurs a t a ratio of peroxide to ferrous salt greatly in excess of that required for the oxidation of ferrous salt to ferric. Discussion h different intermediate, called a ferrous perhydrol, was postulated in 19 13 by C. S.Mummery,” mho wrote the second of a series of papers on LiStudies on Oxidation“ under H. E. . A r m s t r ~ n g 58. ~ ~This ~ improbable compound has assigned to it the structure, Fe(S04H)(OOH). The hypothesis will not explain certain of the properties of induced reactions as determined by Manchot, Goard and Rideal, and others; furthermore his data are easily understood in the light of the higher-oxide hypothesis. Two points only will be considered here. “Vhen Fenton’s agent acts as an oxidizing agent,” writes Mummery, “as it is reconverted into ferrous salt in the process, it is clear that it merely undergoes reduction, exchanging OH for H.” He is so intent upon seeing the depolarizing action in oxidation processes that he shows a tendency to neglect any inconsistency arising in the conclusion. If the intermediate compound were to return to a ferrous salt even momentarily on delivering itself of its oxidizing ability, it would be in a position to react with hydrogen peroxide and to reassume the intermediate form. Thus the ferrous iron xould act in a catalytic manner, which is not at all verified by the factseven those which Mummery himself observed. By “Fenton’s agent” hc apparently means the active intermediate compound formed in a mixture of ferrous sulphate and hydrogen peroxide. Mummery’s error is that he is wrong in his premise, that the intermediate is reconverted to ferrous salt. When a mixture of ferrous sulphate and excess hydrogen peroxide acts as an oxidizing agent, the iron always ends up in the ferric form; thus if it be assumed that the intermediate undergoes reduction, it is reduced to ferric iron, and is therefore reduced f r o m a stage of higher oxidation. I t is obvious that this hypothesis, involving catalysis as it does, fails to rsplain the existence of an induction factor-an equilibrium ratio between acceptor decomposed and ferrous iron oxidized. For, under the hypothesis, a trace of ferrous iron would accelerate the decomposition of any amount of acceptor, and the induction factor would in consequence tend to become infinite. Yet Manchot obtained the induction factor 2 , and the present esperinients show that the limiting factor is 3 . An adequate theory, Mummery mentions, must explain the fact that the “interaction of hydrogen peroxide and ferrous sulphate is one in which initially a ferrous compound is produced which is a powerful oxidizing agent.” But why must it be a ferrous compound:’ Mummery’s reasoning runs along two similar lines, as follows: The active oxidizing agent is either a ferrous or a ferric salt. But the reaction goes quite slowly when ferric salt is used instead of ferrous. (He is speaking of the Oxidation of formic acid.) Hence the intermediate compound is a ferrous salt.

1648

DANFORTH R. HALE

Again, ferrous salt is known to be an oxygen carrier. Hence the oxygen of the hydrogen peroxide is transferred initially to the ferrous salt. But a ferric salt is not formed at once, since the Oxidation of the acceptor proceeds too rapidly. Thus the intermediate is a ferrous salt. These are not cogent proofs; they are hardly probabilities. The possibility of a third form of iron is not mentioned, yet we know of the existence of iron in other forms in the ferrates and perferrates. What everybody admits is that while an induced reaction is in progress, an oxidizing agent more activc than hydrogen peroxide, is present. If ferrous sulphate is causing the hydrogen peroxide to become more actively oxidizing, the reaction belongs under Type I : catalysis with destruction of the catalyst. That this is not the true explanation is shown by the existence of a stoichiometric proportion between the amount of ferrous iron added and the amount of acceptor that may be decomposed-this proportion is, of course, the reciprocal of t,he induction factor. On the contrary if an active oxidizing agent is formed as an intermediate, it is not ordinary ferric iron nor a complex of ferric iron with actor or acceptor, because when soluble trivalent iron compounds are added instead of ferrous iron, the reaction does not follow at all the same course. Two possibilities are left: a complex of ferrous iron with actor or acceptor; a higher stage of oxidation of the iron. Munimery chose the complex with actor, but his is not the correct formula, for it corresponds to Fe0.H202, whereas the present experiments indicate that if the intermediate is a complex with actor its formula is Fe0.zH202. Goard and RideaP in 1 9 2 4 developed a beautiful electrochemical mechanism-hinted at much earlier by Luther and Schilow2?-to explain the course ced reaction when it proceeds through an intermediate higher to account for its becoming catalytic under certain conditions. They also offered further proof of the intermediate formation of iron pentoxide when hydrogen peroxide is the actor. I n its general form an induced reaction proceeding through an intcrmediate higher oxide can be represented as follows: x 2 0 = A02

+

+

+

XOz B = A0 BO SONthe oxidizing ability, as measured with a potentiometer, of A 0 2 will be higher than that of A 0 and BO, and the oxidizing ability of these will be higher than that of A and €3. If the oxidizing ability of B lies below that of A, then, considering an excess of B present, B should be able to reduce XO? not only to A0, but down t o A ; thus A is again in a position to form the peroxide, and will act catalytically. If, on the contrary, the oxidizing ability of B lies above that of A, then B cannot reduce A 0 2 lower than XO, since it cannot reduce a substance to a state lower than its own, and the reaction d l be induced by A . With cerous salts as inductor, Goard and Rideal show that when arsenite is the acceptor, the oxidizing abilities of the chemical substances present assume the induced order, and the reaction is an i n d u e d reaction; and when

I S D V C E D REACTIOSG A S D THE HIGHER OXIDES OF I R O S

1619

a reducing sugar is the acceptor, the oxidizing abilities assume the catalytic order, and the reaction is actually catalytic. In a study of the reaction bctween hydrogen peroxide and potassium iodide induced by ferrous iron, they find that the same relation holds between the character of the reaction and the order of the oxidizing potentials of inductor, acceptor, and oxidized acceptor-in neutral solution the reaction is induced, and in acid solution thc order is altered and the reaction is catalytic. The most interesting part of their work, however, is the development of R new method for the study of induced reactions: the use of an oxidationreduction cell, a calomel electrode, and a potentiometer, for detecting the formation and disappearance of a substance of high oxidizing power in the reaction mixture. Goard and Rideal find that the addition of a little ferrous ammonium sulphate solution to a solution one-thousandth normal with respect to hydrogen peroxide causes a slight rise in potential, which must be due to the formation of another oxidizing agent. This phenomenon is also shown by bichromate and ferrous salt, as has been mentioned. Buffered solutions must be used in order to prevent the superimposed effect of a potential due to change in the hydrogen ion concentration. .Is more of the ferrous solution is added, drop by drop, there comes a time when the potential falls ofE suddenly. The drop is followed by a horizontal section or even a short rise in the titration curve, and this in turn becomes a rather gentle downward slope. The end-point-the lowest part of the sudden drop-occurs at a point corresponding to the destruction of I . j niols of hydrogen peroxide per mol of ferrous sulphate. For the same concentrations of reagents, the general shape of the curve and a break in the curve quite near the ratio o&,; mols hydrogen peroxide per mol ferrous sulphate have been obtained by Lfanchot2‘ and by the writer. -4ccording to Goard and Ridcal the peculiar form of the titration curve indicates “beyond question that the oxidation of ferrous iron bj: hydrogen peroxide proceeds by the formation of an intermediate compound,’’ and since the end-point occurs at this particular ratio, “this intermediate compound must therefore possess the essential forniula Fc20j.“ The relation between this formula and the ratio of 1 : r . s (or 2 : 3 ) may be made clearer by a hypothetical equation: zFeO 3H& = Fe205 3H20. Goard and Rideal iiiiply that the break indicates the complete disappearancc of the hydrogen peroxide, and they conclude that it proves the esistcnce of an intermediate compound of formula I+.Oj in the reaction mixture. AIanchot later verified thc curve, criticized Goard and Iiideal for calling it a “direct proof” of the esivtenee of FesOj, and concluded that the curve is an indirect proof of its cxistence. Goard and Rideal found that, in the neighborhood of the break, the addition of ferrous sulphate to the reaction mixture causes the potential first to drop and then to rise during at least three minutes. Experiments by the writer show that the potential is still slowly rising at the end of the righth minute aftcr a small addition of ferrous sulphatc. Therefore there

+

+

16j o

DANFORTH R . HALE

is occurring a reaction which is producing a strong oxidizing agent. Hydrogen peroxide is the only oxidizing agent originally present, and as its concentration cannot increase, another oxidizing agent is being formed, which must be a higher oxide of iron. Although the minimum in the curve corresponds to the ratio 1.j mols hydrogen peroxide per mol ferrous sulphate, this slow formation of an oxidizing agent continues until much more ferrous sulphate has been added, in fact until the ratio has become 1.14. This formation of a strong oxidizing agent can occur only a t the expense of the hydrogen peroxide; therefore the occurrence of the reaction resulting in the formation of the oxidizing agent is proof that hydrogen peroxide is present’. It follows, then, that hydrogen peroxide exists in the reaction mixture until considerably more ferrons sulphate has been added than corresponds to the formation of an intermediate of formula Fe,Os. Thus Goard and Rideal, and Manchot, are wrong in assuming that the break is simultaneous with the vanishing of the hydrogen peroxide. It is concluded, therefore, that while their curve is evidence for the existence of an intermediate higher oxide, the break in it offers no indication of the formula. The existence of the break and hump is probably to be explained on the basis of the different relative velocities of the disappearance of hydrogen peroxide, and the formation and disappearance of the intermediate, influenced by the varying ferrous and ferric iron concentrations in the reaction mixture. Heinrich IVieland,5gwho was interested in induced reactions from the point of view of biological oxidation processes, published in 1927 an account of a study of the reactions involving ferrous iron and such organic acceptors as oxy-acids, keto-acids, and amino-acids. ‘ T o t solely from the view-point of pure chemistry,” he says, “do we regard the investigation of this subject as desirable. Since we consider that hydrogen peroxide is an intermediate product of biological oxidation, we are at the same time seeking a conception as to what extent this intermediate,product can approach the catalytic action of iron, contained in all cell fluid, in oxidative action in the sense of a peroxydase.” Wieland repeated much of Fenton’s work, determining, however, not the properties of the products of the oxidation, but the relative amounts of organic substance and ferrous iron oxidized. Instead of finding a small ratio between these quantities, as hIanchot, Schilow, and others had found, he observed that the ratio varied up to IO, 20, and even 30, depending on the acceptor employed. Thus he and Fenton observed practically the same thing, that a small amount of iron would accelerate the oxidation of a large non-stoichiometric proportion of acceptor. He admits the possibility of higher peroxides of iron, but as an explanation of the extraordinary ratios he found, he assumes that complex compounds of iron are the cause of the apparent acceleration. He made a quantitative study of the activating power of ferrous iron as a function of hydrogen ion concentration, finding that a t a pH of 3.6 an optimum activation was reached when the acceptor was glycollic acid.

IKDUCED REACTIOSS AND THE HIGHER OXIDES O F IROS

1651

Throughout the investigations Wieland employed solutions so concentrated as greatly to favor the undesirable reactions. The action of hydrogen peroxide on most of the organic acceptors is quite appreciable at the concentrations he used, and the catalytic effect of ferric iron is also large. These reactions] objectionable because they tend to mask the effects of the induced reaction itself, cannot be arrested, but a t great dilution their rates are so low in comparison to the rate of the induced reaction that their effect is entirely negligible, His experiments with potassium iodide as acceptor show clearly the confusing effect of high concentrations of reagents, including acid. Under these conditions the action of ferric iron becomes as great as that of ferrous, and during the first thirty seconds of the reaction, from 6 to 9 equivalents of acceptor are decomposed to one of iron oxidized. The blank run, with no iron added, shows in thirty seconds a decomposition of acceptor which is as large as 119 to 1:’4 of the decomposition effected by the presence of iron, Thus his data are practically valueless for investigating the induced reaction. In answer to this paper of Wieland’s, ManchotGOshortly afterward published a defense of his hypothesis of an intermediate iron pentoxide. He repeated several of the experiments performed in 1901, and by four different methods showed that to each equivalent of ferrous iron oxidized, three equivalents of oxygen disappear into the intermediate, or two are transferred to an acceptor. He repeated Goard and Rideal’s potentiometric experiments with substantially the same results as the originators of the method had obtained; he repeated the induced reaction using potassium iodide as acceptor] confirming the former conclusion; but, unfortunately, he omitted to state how much ferrous iron, or what proportion of ferrous iron to acceptor, was used. His third method is to show that when ferrous iron is added to hydrogen peroxide a t great dilution, the peroxide is detectable in the reaction mixture until in the latter the proportion of reagents is three equivalents of peroxide to one of iron, and then after this point ferrous iron is detectable. His fourth method, a valuable contribution to the technique of studying induced reactions, is to dilute 20 cc. of approximately decinormal hydrogen peroxide to two liters, to add drop by drop four to six cc. of approximately decinormal ferrous sulphate (first generously diluting it) to the strongly stirred solution, and finally to back-titrate the hydrogen peroxide with permanganate. These investigations may be criticized from the lack of attention paid to the extraordinarily rapid reaction between intermediate and ferrous iron, a reaction that undoubtedly occurs to a measurable extent even under the carefully controlled conditions employed in the fourth method. It is noteworthy that although the assumption of Fer% will explain the results by Manchot adequately if the intermediate-ferrous iron reaction is negligible, the assumption of an intermediate of the essential formula Fe03 will explain the results just as adequately, and moreover will not require any assumptions as to the unimportance of this disturbing reaction.

1652

DAKFORTH R. HALE

On increasing the concentration of the hydrogen peroxide solution from about 0.004 normal to 4 normal, Manchot finds that the number of equivalents of hydrogen peroxide disappearing per atom of ferrous iron increases from 3 to 2 4 . 5 . This increase is explained satisfactorily on the hypothesis of a reaction between intermediate and hydrogen peroxide, whereby the latter reduces the former back to the ferrous state. “Obviously it is true,” says Manchot, “that at a high concentration reactions occur which are not occurring noticeably at great dilution.” It has been observed in the case of mangaiiese and lead, that hydrogen peroxide shows the property of first forming peroxides out of them, and then of reducing these peroxides. He continues: “One may even say that this property belongs SO surely to the characteristic properties of hydrogen peroxide, that as soon as the formation of an iron peroxide in the reaction mixture is assumed, this other reaction must be reckoned with.” In proof of the correctness of this hypothesis, Manchot shows that during the reaction in solutions of the greater concentrations, ferrous iron can bc detected in the reaction mixture. .As an indicator he USCP alpha, alphadipyridyl dissolved in acetone, a reagent described by Blsu.61 As a check on this work, ferricyanide is used as indicator. The same results are obtainrd with this reagent; but it is less sensitive and much slower of action. Even in the strong solutions of hydrogen peroxide, the reaction is not catalytic, but comes finally to an end with some hydrogen peroxide unconsumed and all the iron in the ferric and perferric conditions. This effect is easily explained through two reactions: the concentration of hydrogen peroxide is decreasing continuously, thus causing a diminution of the reducing action on the intermediate; and the iriteract,ion of the intermediate with ferrous iron is continuously converting the iron into the inactive ferric condition. As has already been mentioned, Miss Benson”O studied by reaction velocity methods the oxygen-carrying ability of ferrous iron in solutions containing bichromate and sulphuric acids. Her measurements indicate that the oxidation of ferrous iron by chromic acid follows a very different course depending on whether or not potassium iodide is present, and as this fact is cited as evidence against the formation of an intermediate higher oxide of iron, it may pertinently be examined here. I n Miss Benson’s own words: “It has been established in the experimental part of this paper, that, in the absence of iodides, doubling the concentrations of either the acid or the ferrous salt quadruples the rate at which the iron in oxidized; while in their presence doubling the concentration of the acid multiplies it by eight or twelve, and doubling that of the iron, by two. It is hard to see how these facts can be accounted for by a theory which assuiiier that the ‘primary reaction’ is the same in both cases.” Miss Benson implies that if the induced reaction proceeds through the formation of an intermediate higher oxide of iron, the expression for the rate of disappearance of ferrous salt should be the sainc whether an acceptor is present or not. On the contrary this could not be true even in neutral solution, for the presence of an acceptor will influence the interaction of the

I S D U C E D R E . i C T I O S S A S D THE HIGHER OXIDES O F I R O S

1653

interiiiediate compound with ferrous salt, and in acid solution a number of other complications arise. Thus, in the presence of acid, iodides are easily oxidized by air and the catalytic effect of ferric iron is greatly increasrd. Consider the probable situation with regard to the effect of ferrous iron concentration on its own rate of change. Suppose the intermediate to br formed according to the following equation, and suppose the acceptor to react immediately with the product : Fe0

+ z 0 = FeOl

Then if the concentration of oxidizing agent does not change appreciably, dx 'dt

=

-klx

wliere z is the concentration of ferrous iron at time t. This theoretical result is just what Miss Benson obtained experimentally. S o w when no acceptor is present, the intermediate compound must decoinpose practically completely by reacting with ferrous iron. This may be rxpresfed in an equation as follows: 3FeO

+ Fe03 = zFe20a,

and then the rate of disappearance of ferrous iron due to both reactions is given by d ~dt, = - k , ~- k2x3y3., where y is the concentration of intermediate at time t. Miss Benson found experimentally that when iodides are absent, dx,'dt = -k3x?, but it is likely that, under certain experimental conditions and within experimental error, the expressions, (-k3s2) and ( - l w - k2x3y),may be numerically indistinguishable. If the above chemical equations are not accurate, (if, for example, Fe0.zH202should be written in place of FeuL)a very similar argument holds. Thus the experimental fact that the rate of oxidation of ferrous salt changes when a suitable acceptor is added, does not preclude the possibility of the primary formation of a higher oxide of iron; but is, on the contrary, just, what would be expected if an intermediate higher oxide were formed.

Permanganate as Actor The reaction involving potassium permanganate and :odium orthoarsenite was investigated under carefully controlled conditions, and was found, likewise, to yield an induction factor of three. So far as the writer is aware, this induced reaction has never before been investigated nor even mentioned. Its main point of interest here is that it is another induced reaction occurring in a non-acid medium, and therefore its course is not complicated by the undesirable retardation and catalytic effects which are brought into play by the presence of acid. I t was necessary to run the reaction in an alkaline solution; otherwise the arsenite would undergo oxidation by t,he air. Soqum bicarbonate was uspd as the alkali. Since ferrous iron is quickly oxidized by oxygen in alka-

I654

DANFORTH R. HALE

line solution, the reactibn had to be run in the absence of air. For this purpose a filtering flask was used as the reaction vessel, and carbon dioxide was allowed to bubble through the solution and to pass out through the side-arm of the flask. Since potassium permanganate in alkaline solution is ordinarily reduced by ferrous iron only as far as the insoluble manganese dioxide, which obscures the end-point, an excess of solid sodium pyrophosphate was added. This reagent increases greatly the reducing ability of the ferrous iron, and the permanganate then is reduced to a soluble manganous salt. The filtering flask was fitted with a three-hole rubber stopper through which passed the tip of a I O cc. burette, the stem of a funnel provided with a glass stopcock near the stopper, and a delivery tube reaching nearly to the bottom of the flask. The other end of the delivery tube was connected by a length of rubber tubing to a cylinder of carbon dioxide. This is essentially the apparatus with which Job62titrated iron in alkaline solution with permanganate. The arsenite solution was added each time directly to the clean flask. The air was washed out by inserting the stopper and passing a stream of carbon dioxide through the flask for a few minutes. To remove the air in the stem of the funnel, the latter was partly filled with distilled water, and the gas was made to bubble up through this by stoppering the side-arm of the flask with a finger and opening the stopcock for a few moments. The stop-cock must be closed again as soon as the gas is released through the sidearm; otherwise the distilled water at once drops into the flask and refills the stem with air. After washing the air out of the flask and out of the funnel stem, the ferrous solution may be added through the funnel or the permanganate through the burette, with the assurance that any oxidation which may occur is due solely to the permanganate. Care must be taken, of course, to wash down all the solution in the funnel without washing in any bubbles of air. The induced reaction was run as follows: Fifty cc. of approximately decinormal arsenite solution and about three grams of pulverized pyrophosphate were put into the flask, and the latter thoroughly washed out with carbon dioxide. Then permanganate was added until the solution assumed a slight pink color, and this volume was taken as a blank correction. (The color would persist overnight with no noticeable change of tint, showing that under these conditions permanganate does not oxidize arsenite.) The correction volume being found, a given amount of ferrous solution, measured from a graduated pipette, was added slowly through the funnel. During this addition, permanganate was added drop by drop, but not in such quantity as t o bring back the color of the blank run. The object of this procedure was to keep the ferrous iron concentration always as low as possible, thus reducing to a minimum the interaction of ferrous iron and intermediate. When the last of the iron had been added, the funnel was washed with three or four portions of water, and then with the permanganate the color of the solution was brought to that of the blank run as nearly as possible. From

IXDUCED REACTIOSS AND THE HIGHER OXIDES O F IROS

1655

the amount of iron added and the corrected volume of permanganate, the quantity of arsenite oxidized and the induction factor were calculated. In Table I1 all numbers, except the induction factors, are volumes in cc. TmLE

Ferrous Iron

11

Potassium Permanganate n

n=o.oj

n=I

blank

4.0

.2

.65

5.15

3.0

.I5

4.2

I .j

,075

1.0

.05

.7 .58 .57

titre

2'35 1.90 1.8

corr.

=

0.158

n=x

Arsenate = perman.

Induction Factor

-iron

4.50 3.5

'71 .55

.5'

2'55

.40

2.67

1'77

.28

.zoj

2.73

1'33 1.23

.21

.16

3.2

'I94

,144

2.88

The volumes in the column headed n = I , are obtained by multiplying together the volumes of ferrous iron or the corrected volumes of permanganate and the normality of these solutions. In the experiment with I cc. of iron, it was impossible to tell whether the permanganate titre was 1.8 or 1.9 cc. That is, with 1.8cc. added, the solution seemed to be of the same color that the blank run had been, and after two more drops of permanganate had been added (making the titre 1.9) the color still seemed the same. Two more drops definitely made the reaction mixture darker in color. The mean of the two values is thus probably more correct than either, and this corresponds to an induction factor of 3.0. These experiments therefore confirm the conclusion previously stated, that in the limiting, ideal case, ferrous iron possesses the ability to transfer three equivalents of oxygen to a suitable acceptor, and hence the intermediate compound may be of the formula Fe03. Furthermore the result makes it very improbable that the intermediate compound is an addition product of actor and inductor, since, if this were the explanation, the intermediate for hydrogen peroxide would be Fe0.zH20?,and for permanganate, .;Fe0.4KMnO4. Oxygen as Actor H islor ical After discovering that Manchot's assumed pentoxide of iron's, Fe205,is actually essentially a trioxide, Fe03, it was natural to hope that experimental error had led him to the conclusion that a dioxide of iron is the intermediate when oxygen is actor, and that this case, too, would turn out to be intelligible on the assumption of an intermediate of formula FeOa. Then whether hypochlorous acid, oxygen, hydrogen peroxide, or potassium permanganate were the actor, the reaction would follow the same general course: it would proceed through the intermediate formation of an unstable compound representable as FeOB,and iron would present one uniform behavior toward these oxidizing agents instead of an improbable variety. But this simple state of affairs does not exist. In the presence of ferrous iron, hydrogen peroxide will liberate iodine from potassium iodide, but at-

16 j h

DASF’OHTH R. I I h L E

niosplieric osygm under ordinary pressure will not. This was shown in an wperinient performed as a blank in connection with the study of hydrogen pcroxide as actor: a dilute solution of potassium iodide and ferrous iron was thoroughly aerated by rotating the -tirrer at a high speed, but no iodine w a , ~ set frce. Hence thc assumption of only one intermediate iron compound will not explain the action of both hydrogen peroxide and oxygen. If FeOI w r c fornied in each caw, then in each case iodine ought to be liberated. The manner in which JIanchot investigated the oxygen-arsenite reaction ha. alrpady hcen described briefly. I t consisted in measuring the oxygen absorbed nhcn a solution of pota,ssiuni orthoarscnite and ferrous sulphate arp shaken in an apparatus filled with the gas. The color of such a mixture is first a light green, and, according to llanchot, “the red-brown color of ferric hydroside first appears w h m there i3 consumed almost double the volunie of osygen corresponding to the change froni ferrous to ferric iron.“ Typical experiments were perforiiied with the following proportions of reagents: I.; g. ar.srnious oside, 36 p. potassium hydroside, I O O cc. water, and 1 0 or 15 cc. approximately tlccinormal ferrous sulphate in 30 cc. of water. It was found that the absorption of oxygen continued after the iron had been osidized-although at a much slower rate. \\*hen the observed volume of osygen absorbed was corrected as carefully as possible for the effect of this side reaction, it conipared with the calculated volume (twice that required in the osidation of the iron alond as shown in Table 111. ‘TABLE

I’errous sulphate 10

15

cc

111

Observed vol.

.6cc. 37.7 2 j

Calculated vol 2j

.6cc.

37.6

Practically the identical erperiinents were run in 192 j by James H. as an atom. As further evidence for the widespread nature of this action of oxygen, von Baeyer and l-il1igerGsshowed that when benzaldehyde is oxidized in air, the oxygen combines as the molecule, forming benzoyl hydrogen peroxide, which then reacts with an acceptor (more benzaldehyde, indigo, etc) losing an atom of oxygen and yielding benzoic acid. In their study of the oxidation of ferrous iron by osygen, Smith and Spoehr6’ obtain reaction velocity data for a first order reaction: the halfperiod of the reaction remains constant during a three hundred percent increase in the proportion of ferrous salt added. They suggest that the reaction is a case of so-called “autoxidation,” that the iron unites with molecular oxygen to form a “moloxide,” and that this intermediate reacts with the acceptor if one is present. They make the further coininent, “It may be of significance in this connection that stoichionietrically the moloxide, FcO-O!, is equivalent to a ferrate.*” They assume that the ferrous atoms combine singly with the oxygen, and that on this account and since the oxygen pressure was made constant, the reaction appears to be monomolecular. Following this line of thought’, t,he simple molecule, FeO-Oa, would be the most likely substance to be formed. Rut if this were true, the volume of oxygen absorbed would not be twice the volume equivalent to the ferrous iron, but four times this volume. Since, however, the experimental data show the former value-twice the volumethe empirical formula of the intermediate inay be FcO?. The relations between the empirical formulae and the volumes of absorbed oxygen inay be made clearer through hypothetical equations:

+ O2 = Fes04 Fe?O, + 0 + no2 z(FeO-02) = Fe203+ 3 0

Fe202 FcaOs

=

=

In the first equation the equivalent of iron reacts with I ‘ 2 02,which is just double that reqiired to osidize the equivalent of iron up to the ferric state, namely, 1 / 4 O?. In the second equation the equivalent of iron reacts with O?, which is four times that required to oxidizc the equivalent of iron. Thus the experiments show that the second equation does not esplain the phenomenon. This conclusion makes an examination of the first equation desirable. 0,for the ferrous salt is reacting with The reaction cannot be written: FeO molecular oxygen; and the reaction is obviously not niononiolecular if it is written: nFeO O?. However, if the sodium ferro-pyrophosphate be considered to possess the formula proposed by Pascal,71SasFcr(P20ija,then the

+

+

* The term “moloxide” is explained by the contest as an oxide in which oxygen enters as the whole molecule. TVritten “mol-oxide” it is used by \Yelo and B a u d i ~ c has~ ~an equivalent for the older term, “holoxide,” invented by Traube.’O “Ferrate” is the older term for a salt having hexavalent iron in the acid racl&al; in the present thesis, folloiving J. S . Friend,4l the term for such a salt is “perferrate, and “ferrate” is kept for the compounds containing quadrivalent iron.

ISDUCED REACTIOSS .4SD THE HIGHER OXIDES O F IROS

1661

reaction can be monomolecular with respect to the iron compound, for the latter can react with molecular oxygen, and the intermediate higher oxide can reasonably be written FezOt. Actually the reaction cannot be monomolecular even with respect to the ferrous iron, for the higher oxide can return to the ferric stage only by interaction with more ferrous iron; thus the rate of disappearance of thc ferrous iron cannot be proportional to the concentration of the ferrous iron alone; it must be proportional to a term involving also the concentration of intermediate. I t has been a stumbling block in the theory of induced reactions which postulates intermediate higher oxides that, when a ferrous salt alone is oxidized in dilute solution with oxygen or hydrogen peroxide or certain other oxidizing agents, only so much of the oxidizing agent is consumed as corresponds stoichiometrically to the oxidation of the iron from the ferrous to the ferric condition. If an intermediate higher oxide be first formed, and even if it disappear through reaction with unused ferrous iron to form ferric iron, some of the iron at the end of the reaction ought still to be in the highly oxidized state; and this amount would correspond to a quantity of oxidizing agent beyond that theoretically required, Yet an error of this sort, if it does occur, is too small to be detected by the ordinary analytical method. The data of Smith and Spoehr, mentioned above, show this. The oxygen absorption calculated from the equation, qFe++ O2 = qFe+++ 2 0 = , was 2 7 . 4 cc. a t atmospheric pressure for I g ferrous sulphate heptahydrate, and the observed volumes for two experiments were 27.5 and 2 7 . 6 cc. That the reaction is experimentally monomolecular (when the pressure of oxygen is constant) and that, within the limit of analytical accuracy, no over-consumption of oxygen occurs, permits the deduction of two important facts: I ) Ferrous iron reacts two atoms a t a time in this oxidation; 2 ) The interaction of the intermediate higher oxide and ferrous iron is remarkably rapid. Consider the fundamental meaning of "monomolecular" as applied to a chemical reaction. When Smith and Spoehr63 say that the reaction is monomolecular, they mean actually that the half-period of the reaction is a constant, which, as is well known, is true only if the rate of change in concentration of one of the reactants or of the only reactant is directly proportional to the concentration of this reactant. Since the simplest case of this proportionality relation is one in which single molecules decompose, the relation is often loosely taken as a definition of the monomolecular reaction. There are other cases, however, involving not merely single molecules, in which the rate of change of a substance is directly proportional to the concentration of that substance. Any explanation or reaction mechanism fitting such cases must show how it happens that the rate of change is proportional to the concentration of the one component only. Of course under ordinary circumstances the rate actually is proportional to a more cpniplicated expression, but because of their lack of delicacy, analytical methods yield data which fit the complicated expression and the simple proportion equally well. Thus

+

+

1662

DANFORTH R. HALE

when the data for the hydrolysis of an ester point to a monomolecular reaction, they are satisfying the integrated monomolecular equation, dCl/dt = - k G , as well as they wduld satisfy the integrated form of the reaction velocity equation that expresses the true situation, namely, dCl/dt = - k’C1C2,where C1 is the concentration of ester and Cs is the concentration of water (mols of water per liter of solution) at time t . The essential thing here is that C z is so nearly constant that for all practical purposes, and for all ordinary methods of analysis, the product of IC‘ and Cz maybe regarded as being a new constant. A hypothetical reaction to explain the oxidation of ferrous salts may be written in two steps as follows:

+ O2 = Fe20J Fe2O2+ FezOl = zFe203 Fe202

At time t let z be the concentration in mols per liter of the reactant containing ferrous iron, and let y be the corresponding concentration of intermediate. Then from the first equation the rate of disappearance of the ferrous compound is given by dx/dt = - klx”, where n represents the number of ferrous molecules combining a t a time, and from the second equation, dx/dt = -k2xny. Then the total rate of change of the ferrous salt is dx/dt = -klx” - k2xny. Since the reaction has been shown to be apparently monomolecular, the rate of change of the ferrous salt must be of the form dx/dt = -kx. This can happen only when n = I , and when the term -k2xny is very small compared to the term - klx. I n order for n to be I , the ferrous iron must enter into the reaction two atoms at a time, which may be accounted for by assuming the equilibrium zFeO = FezOz, or by assuming that the actual reactant is Pascal’s pyrophosphate. In order for - k2xnyto be very small in comparison to - klx, k2 must be verysmall in comparison to kl,or y must be very small. Now klis a measure of the rate of disappearance. Since no over-consumption of oxygen occurs in this oxidation, the intermediate obviously does not pile up in the solution; hence the rate of disappearance of the intermediate is greater, probably much greater, than its rate of appearance; thus kz is greater than kl. Therefore y is a very small quantity, probably a t least a thousand times smaller than z. Thus may be explained how it happens that the oxidation as a whole follows the velocity relation of the monomolecular reaction. Conclusions I. The induced reaction in which ferrous iron plays the part of inductor belongs to Type 3 of Bancroft’s classification, the acceptor reacting with a higher stage of the inductor. 2. When ferrous iron is acted upon in neutral solution by hydrogen peroxide, or in alkaline solution by potassium permanganate, the intermediate compound formed has the essential formula Fe03. 3. When ferrous iron is acted upon in alkaline solution by oxygen, the intermediate compound has the essential formula Fe20a.

I S D V C E D RE4CTIORS A S D THE HIGHER OXIDES O F IRON

1663

4. I n an induced reaction involving ferrous iron as the inductor, the acceptor is decomposed because the intermediate higher oxide is a more powerful oxidizing agent than the actor. j. The difference between the action of hydrogen peroxide and potassium permanganate, and oxygen, must be ascribed to the electrochemical potentials of peroxide and permanganate being greater than that of oxygen. 6 . Manchot’s discovery that different oxidizing agents yield different intermediate compounds is thus confirmed. That he obtained the wrong formula, FesOj for the compound when hydrogen peroxide or potamium permanganate is actor was due to a fortuitous choice of concentrations of the reagents. 7. From potentiometric data Goard and Rideal concluded erroneously that the intermediate had the formula Fe205. The conclusion was based on the incorrect assumption that a sudden drop in their curve indicated the disappearance of the hydrogen peroxide. Actually this portion of the curve yields no hint of what the intermediate higher oxide may be.

Acknowledgment Professor Wilder D. Bancroft has been the inspiration for this experimental research as well as the director of it, Feeling that association with him during the past four years has greatly helped the writer to appreciate the meaning and the method of chemical research, the latter gratefully thanks him for his helpful counsel and generous assistance.

BIBLIOGRAPHY Schonbein: J. prakt. Chem., 75, 78 (1858). Mohr: “Titrirmethoden,” 217 (18jj). Kessler: Pogg. rlnn., 119, 218 (1863). Hedges: J. Chem. Soc., 969 (1928). 5 Milas: J. Am. Chem. SOC.,49, zoo5 (1927). 5 Ostwald: Z. physik. Chem., 34, 248 (1900). Mellor: “Chemical Statics and Dynamics,” 333 (1904). Kagner: Z.physik. Chem., 28, 33 (1899). Dhar: J. Chem. SOC.,111, 690 (1917). Schonbein: J. prakt. Chem., 74, 328 (1858). ” v a n ’ t Hoff: Z. physik. Chem., 16, 411 (1895). 1, Jorissen: Ber., 29, 1 7 0 7 (1896). Engler and Keissberg: “Iiritische Studien iiber die .~utoxydationsvorgtinge”(1903). ‘4 Schonbein: J. prakt. Chem., 79,66 (1860). Brodie: Phil. Trans., 141, 759 (1850). l 6 ?+lellor:“Chemical Statics and Dynamics,” 304 ff. (1904). li Traube: Ber., 15, 666 (1882). Bach: Compt. rend. 124, 9j1 (1897). Manchot: ;Inn., 325, 93 (1902). hfanchot: Ann., 314, 177 (1899). 21 Manchot: Ann., 460, 179 (1927). ?%Lutherand Schilow: Z. physik. Chem., 46, 777 (1903). 23 Miller: J. Phys. Chem., 11, 9 (1907). 24 Benson: J. Phys. Chem., 7, 356 (1903).

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Bancroft: J . Phys. Chem., 33, 1184 (1929). Vannoy: J. Phys. Chem., 33, 1593 (1929). 2 i Keir: Phil. Trans., 80,359 (1780). 2d Faraday: Phil. Mag., 9, 122 (1836). l 5 Bennett and Burnham: J. Phys. Chem., 21, 107 (191;). 3 @ E v a n sJ. : Chem. SOC.,131, 1020 (1927). 31 Freundlich, Patscheke, and Zocher: Z. physik. Chem., 130, 289 (1927). 32 Forbes and Bartlett: J. Am. Chem. SOC., 35, 1527 (1913). 33 Furman: J. Am. Chem. SOC.,50, 273 (1928). 34 Baudisch and Welo: J. Biol. Chem., 61, 261 (1924). 35 Dhar: J. Phys. Chem., 28, 943 (1924). 36 Fenton: J. Chem. Soc., 65,900 (1894). 3 7 Fenton: Proc. Chem. SOC.,14, 119 (1898). 3 8 Fenton: J. Chem. SOC., 77, 69 (1900). R a l t o n and Christensen: J. Am. Chem. SOC.,48,2083 (1926). 4 @ Manchot: Z. anorg. Chem., 27,420 (1901). 41 Friend: “Textbook of Inorganic Chemistry,” Vol. IS, P t . 11, 130 (1921). 4 2 Fremy: Compt. rend., 14, 442 (1842). 4 3 Smith: Phil. Mag., 19, 302 (1841). 4 4 Rose: Pogg. Ann., 59, 321 (1843). 4 5 de Mollins: Ber., 4, 626 (1871). 4 6 Moeser: J. prakt. Chem., (2) 56, 425 (1897). 4 7 Retgers: Z. physik. Chem., 10, 529 (1892). l 8 Baachieri: Gaze., 3611, 282 (1906). 4 8 hloeser and Borck: Ber., 42, 4279 (1909). Pellini and Meneghini: Z. anorg. Chem., 62,203 (1909). Goralevich: J. Russ. Phys. Chem. SOC.,58, 1129 (1926); Chem. Abs., 22, 1294 (1928). Williams: “Cyanogen Compounds,” I68 (1915). j3 Goard and Rideal: Proc. Roy. SOC.,105.4, 148 (1924). j4 Mummery: J. SOC. Chem. Ind., 32, 889 (1913). j5 FTeltsien: Ann., 138, 129 (1866). j6 Manchot: Ber., 34, 2479 (1901). Armstrong and Colgate: J. SOC.Chem. Ind., 32, 391 (1913). 3 8 Colgate: J. SOC.Chem. Ind., 32, 893 (1913). 5 9 Wieland: Ann., 457, I (1927). 6o Manchot: Ann., 460, 179 (1927). 6‘ Blau: Monatsheft, 19, 647 (1898). 62 Job: Ann. Chim. Phys., (7) 20, 205 (1900). 63 Smith and Spoehr: J. Am. Chem. SOC.,48, IO; (1926). 64 Just: Z. physik. Chem., 63, 385 (1908). 65 Friend: “Textbook of Inorganic Chemistry,” Vol. I X , P t . 11, 191 (1921). 66 Baur: Z. anorg. Chem., 30, 251 (1902). 6 7 Engler: Ber., 36, 2642 (1903). b8 von Baeyer and Villiger: Ber., 33, 1569 (1900). 65 Welo and Baudisch: J. Biol. Chem., 65, 2 1 j (1925). io Traube: Ber., 26, 1481 (1893). Pascal: Ann. Chim., (8) 16, 386 (1909). 25

26

Cornell University.