THE MECHANISM OF IRON CATALYSIS IN CERTAIN OXIDATIONS

BY C. V. SMYTHE*

(From the Laboratories of The Rockefeller Institute for Medical Research, New York)

(Received for publication, November 8, 1930)

It is well known t,hat ferrous iron is comparatively stable in acid solut,ion and that it is rapidly oxidized to the ferric state by the oxygen of the air in alkaline solution. No satisfactory explanation for this has been brought forward. One difference in the condi- tion of the iron in the two solutions immediately suggests itself. This is that in the acid solution the iron exists as ferrous ion while in the alkaline solution it exists as unionized ferrous hydroxide. If we remember that the oxidation of ferrous iron consists of the loss of one electron from the kernel of the iron atom, we might expect that ionized and unionized ferrous iron would differ in the ease with which they may be oxidized. Thus, the oxidation of ferrous ion would involve the separation of a negative charge from a kernel that is already carrying two positive charges. This would be comparatively difficult. On the other hand, unionized ferrous hydroxide should be more readily oxidized for here it is only neces- sary to separate an electron from a previously electrically neutral substance. With this consideration to guide us we would expect that if ferrous iron can be obtained in an unionized compound it will be readily oxidized regardless of the acidity of the solution.

In a previous communication Smythe and Schmidt (1) have reported studies on the compounds which form very slightly ionized complexes with ferric ion. If the suggestions developed by them concerning the manner in which these complexes are formed are correct, then these same compounds should form fer- rous complexes and, in accordance with the above considerations, the oxidation of the iron contained in the complexes should depend

* National Research Council Fellow in Biochemistry. 251

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252 Mechanism of Iron Catalysis

upon the acidity of the solution only in so far as the stability of the complex depends upon this acidity. We would expect then that in any given solution the rate of oxidation would be propor- tional to the concentration of the ferrous complex. The concen- tration of this complex is determined by the concentrations of the ferrous ions and the complex-forming anions. The concentration of the anions is, within a certain pH range, determined by the pH of the solution; so we would expect the rate of oxidation within this pH range to vary as the pH is varied.

130

30 50 70 90 110 130 Minutes

FIQ. 1. Each flask contained 0.8 cc. of 0.1 M pyrophosphate solution and 0.2 cc. of 0.1 M FeSOd. The O2 equivalent of the iron present is 120 wnm.

An excellent substance to test this hypothesis is pyrophosphoric acid. This gives a ferrous complex which is stable over a wide pH range, and the ferrous iron is the only substance in the solution that can be oxidized. Experiments at various pH values have been carried out with this acid and the results are shown in Fig. 1. It is readily seen that they bear out the above predictions.

The experiments were carried out in the Warburg respiration apparatus. Brodie’s solution was used as manometer fluid. The temperature was maintained at 25”. The total volume of the flask was about 25 cc. and 1 cc. of liquid was used. To get reproducible results it is necessary to use the same or very similar vessels. Ves-

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C. V. Smythe 253

sels of different shape allow different degrees of shaking and this affects the rate of oxidation. Only a few minutes are required for the solutions to come to equilibrium before mixing. A similar flask without any iron was run as a control thermobarometer.

The pyrophosphoric acid solutions were made by dissolving cryst,als of Na4P207e10 Hz0 in water and adjusting the pH by the addition of hydrochloric acid. The pH values recorded were determined by the hydrogen electrode and refer to the pyrophos- phate solution before the addition of the ferrous sulfate. The

110

50 70 90 110 130 Minutes

FIG. 2. Each flask contained 0.8 cc. of 0.2 M metaphosphoric acid solution and 0.2 cc. of 0.1 M FeS04. The 02 equivalent of the iron present is 120 c.mm.

addition of the iron salt would change the pH some, but the solu- tions are certainly in the correct order.

Experiments with metaphosphoric acid have been carried out in t,he same way and the results are shown in Fig. 2. In general the results are the same as those with pyrophosphoric acid, but it may be seen that the rate of oxidation at any one pH is not the same in the two cases. This is to be expected from the above considerations since the structure and the dissociation constants of the two acids are not the same. The metaphosphoric acid solutions were made by dissolving the free acid in water and adjust- ing the pH by the addition of sodium hydroxide. When ferrous sulfate was brought to pH 2.0 by means of sulfuric acid instead of

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254 Mechanism of Iron Catalysis

metaphosphoric acid, no oxidation was observed within a compar- able length of time.

Perhaps the best known ferrous complex is the ferrocyanide. The iron in this complex is certainly not ferrous ion and yet its oxidation by the oxygen of the air is exceedingly slow (2, 3). The methods used by the above authors in studying this are open to severe objection and this compound needs special investigation. In any case, in line with the reasoning of this paper, one would expect that the iron in this compound would not be readily oxi- dized, for if we consider the structure of the ferrocyanide ion to be

CN CN = [ 1 CN Fe CN CN CN

we can see that the iron atom is surrounded by four negative charges. In order for an electron to escape from the iron it must penetrate this negative atmosphere and this is not easy to do. If one could repress the ionization of these cyanide groups so they would exist as CNH instead of as CN-, it should then be easier for the iron to be oxidized. This repression is very difficult, how- ever, for ferrocyanic acid is a strong acid. Kolthoff (3) has esti- mated that its fourth constant is 5 x 10-4. The third one was too strong for him to estimate. It is possible to reduce this negative charge by substituting an NH, group for one of the CN- groups. The iron of this substance, pentacyanoammine-ferroate, is oxidized by molecular oxygen at an appreciable rate (4). Baudisch and Davidson (5) found the oxygen consumption to be much more rapid at a pH of 2.0 than at pH 7.0 or 12.0. It is striking that here increase of pH slows down the rate of oxidation in contrast to the former cases. They interpreted this in terms of the following equation, H+ + 30, (disso lved) + 2[Fe(CN)5.NHI]- *OH- + 2[%e(CN)5NH3]=, which shows the hydrogen ion taking part in the reaction. In accord with our other considerations we would regard the influence of the increased hydrogen ion concentration as due to a repression of the ionization and consequent decrease of the nega- tive charge on the iron compound.

We know that the hydroxy organic acids are good complex formers (1) so t,hey should unite with ferrous iron and make it readily oxidizable. The results obtained with these acids are more complicated, however, for the organic radical is also oxidized

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C. V. Smythe 255

and we are led into a consideration of catalysis by iron. Many people have worked upon this problem and no attempt will be made to review all of the results. We shall confine our discussion to a few papers closely related to our immediate problem. Wieland and Franke (6) have published an exhaustive paper upon the oxida- tion of organic acids by molecular oxygen in the presence of ferrous salts. In other papers (7, 8) they have used hydrogen peroxide instead of molecular oxygen as the oxidant. Their general conclu- sion is that the ferrous ion unites with the substrate thereby mak- ing the latter more susceptible to oxidation. Manchot (9) and others (10, 11) prefer to believe that the ferrous iron forms a peroxide and that this peroxide is the oxidant. One can recognize here the outlines of the two general theories of oxidation, that of peroxide formation and that of hydrogen activation,

If we apply to this problem the ideas stated at the beginning of this paper and others which follow from it we are led to the follow- ing considerations. If ferrous sulfate is added to tartaric acid, for example, some ferrous tartrate is formed, the amount formed depending upon the pH of the solution. We may assign to it the formula of Substance 1 in the accompanying scheme.

Substance 1 Substance 2

e -- -2H+ -+

Substance 3

2e --+

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Substance 5

0 Substance 6

Substance 4

//O C

\ \

\ HC-0 0

Rearrange 3 i ‘>e+ - 2 H++

\o.f.‘. ( 0

1’

/

\ 0

Substance 9

Substance 7

No C

\ \

\ c-0.. 0

II I

5.. I “Fe+

,.*’ I

c-o**‘. 0

I/

/

cl 0,

‘0

Substance 10

No c\ O, I \ + co2 \

Substance 8 C=O Fe+ --+

Ho O/

o,Fe++ co*

C / /

/ C

I\ / /C

c=o 0 I

\O ”

Rearrange -* Fe+

1 % R0

c=o 0 40

IV b# c\ 0

$02

c\

0, / C-OH

Fe+ -I- 1 0 C-OH

0’ \

C’ 0

256 Substance 10

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C. V. Smythe 257

The dotted bond line is used here in the same sense that it was used by Smythe and Schmidt, to represent an attraction due to the residual negative charge on these oxygen atoms. For the calcula- tion of this residual charge the reader is referred to the previous paper (1) or to the paper by Latimer and Porter (12). This iron is now electrically neutral ferrous iron so it can readily lose an elec- tron to an acceptor such as molecular oxygen. The loss of this electron gives the iron a positive charge so it will draw in the electrons represented by the dotted lines. This makes the oxygen atoms involved more positive, in other words it makes the OH groups more acidic, and the hydrogens attached to them are to an appreciable extent forced off as hydrogen ions. In order to avoid confusion later we may point out that although we have ferric iron as soon as the electron separates from the ferrous iron our compound is not the usual ferric tartrate. The usual com- pound, formed when preformed ferric ion is added to tartaric acid, has the three iron valences attached to three carboxyl groups.

The electronic arrangement of the oxygens may now be repre- sented as follows:

The iron is attracting a pair of electrons from each of the two oxygens.’ If we consider these electrons to be in a dynamic con- dition it is easy to see that at some instant the iron will attract one pair to a greater extent than the other. This makes one

1 This may be regarded as what, in older terminology, was called a “eersplitterte” valence (13).

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258 Mechanism of Iron Catalysis

oxygen more positive than the other and the pair of electrons least attracted by the iron will be drawn toward the more positive oxy- gen, thus forming a bond between the two oxygens. The other pair of electrons passes to the iron and is immediately lost to an acceptor (molecular oxygen); or, stated in other words, the ferric ir6n is reduced and immediately reoxidized to the ferric state by the oxygen. Our compound is now represented as follows:

No

This ring involving the oxygen-oxygen linkage will rearrange to form Substance 5 in our scheme for this rearrangement is accom- panied by a decrease in the sum of the products of the kernel charges of the atoms attached together (14). We are now in a position to repeat the process and lose two more hydrogens forming Subst.ance 8. This will break down to form either 2 molecules of oxalate or it will lose 1 molecule of CO2 at a time, first forming Substance 9. That the process actually occurs in a manner very similar to the above may be shown by the following facts. Fenton (15), who was one of the first to study this reaction obtained dihydroxymaleic acid (Substance 5 in the above scheme) in crys- talline form from a mixture of tartaric acid and ferrous iron. He believed that this lost COz-forming Substance 9 in our scheme; but Wieland and Franke from a study of the 0, consumed and COZ produced as well as by identification of the compounds themselves showed that the main reaction was 60 form Substance 8, which then gave Substances 9 and 10 in the ratio of 2: 1. They found that Substance 9 slowly evolved COZ. The fate of the ferric oxa- late, Substance 10, will be discussed later.

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C. V. Smythe

If instead of adding ferrous iron to the tartrate we had added ferric iron the results would not be the same for here there is nothing to start the process; i.e., there is no easily detachable electron analogous to the one possessed by ferrous iron. It does not follow, however, that ferric iron has no effect on the oxidiz- ability of the tartrate. It should make it definitely more sus- ceptible to oxidation for its secondary valences neutralize the residual negative charges on the oxygens of the alcoholic hydroxyl groups thereby making it easier for the hydrogens of these groups to separate.

Again the available evidence supports these predictions. Both Fenton (15) and Wieland and Franke (6) found that whereas ferrous ion markedly catalyzes the oxidation of tartaric acid, ferric ion is without measurable effect. However, when Wieland and Franke used a less inert oxidant, hydrogen peroxide, ferric ion showed a slow but definite catalytic effect. Similarly when they used a more readily oxidizable substrate, dihydroxymaleic acid, its oxidation by oxygen was catalyzed by ferric ion. This is to be expected for the alcoholic hydroxyl groups of this acid are more strongly acidic than those of t.artaric acid, due to the adjacent double bond, and the secondary valences of the ferric iron enables them to ionize. In fact, a careful perusal of Wieland and Franke’s entire paper has revealed nothing inconsistent with the considera- tions advanced here. The marked change, at higher pH values, in their curves for the oxidation of tartaric acid and dihydroxymaleic acid may require some explanation. I have repeated their experi- ments with tartaric acid, with solutions made up as they describe, but with only 1.6 cc. instead of the 20 cc. that they used, and I have obtained essentially the same results as Fig. 3 shows. The pH values recorded here are those reported by Wieland and Franke and refer to the solution after the ferrous sulfate has been added.

The striking thing about the above results is that the oxidation drops off very suddenly in the more alkaline solutions. If we remember that in these solutions we are dealing with a competition between tartrate and hydroxyl ions for the ferrous or ferric ions as the case may be, the explanation will become apparent. When the tartrate and ferrous sulfate are mixed, the higher the alkalinity of the solution the more ferrous hydroxide is formed. We know that ferrous hydroxide is rapidly oxidized and we know t,hat the ferric

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260 Mechanism of Iron Catalysis

70 90 110 130 150 160 Minutes

FIG. 3. Each Aask contained 1.52 cc. of 1.0 M tartaric acid solution and 0.08 cc. of 0.5 M FeS04. The O2 equivalent of the iron present is 240 c.mm.

220, I

MinuTes

FIG. 4. Each flask contained 0.8 cc. of 0.2 M citric acid solution and 0.2 cc. of 0.1 M FeSO4. The 02 equivalent of the iron present is 120 c.mm.

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C. V. Smythe 261

Minutes

FIG. 5. Each flask contained 0.8 cc. of 0.2 M oxalic acid solution and 0.2 cc. of 0.1 M FeSOa. The 02 equivalent of the iron present is 120 c.mm.

24 $i 70

& k 50 s

0. Iii 30

IO

10 30 50 70 90 110 130 150

pH 2.9

-ib---- Minutes

FIG. 6. Each flask contained 0.8 cc. of 0.2 M oxalic acid solution and 0.2 cc. of 0.1 M FeSOa.

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262 Mechanism of Iron Catalysis

iron so formed, even though it combines with tartrate, will have no catalytic action. Then if all the iron formed ferrous hydroxide the solution should take up just enough oxygen to oxidize the iron to the ferric condition and no more. Apparently Wieland and Franke obtained just this result at pH 9.0. The results shown in Fig. 3 do not show such a sharp end-point. Another factor that contributes to these results is the format,ion of ferric hydroxide from the ferric tartrate complex. We know that ferric hydroxide is a very insoluble substance and would tend to form to some extent at the more alkaline reactions.

Fig. 4 shows that the oxidation of citric acid at various pH values occurs just as our theory would predict. The pH values recorded refer to the citric acid solution before the addition of ferrous sul- fate. It must not be imagined that the details given for the course of the oxidation with tartaric acid will apply without modification to the other acids.

We may now discuss the results with the simpler acid, oxalic, which may be an intermediate product in the oxidation of the above acids. The only oxidation product of this acid is carbon dioxide, so its determination is of interest to us here. It may readily be determined by running two vessels, one containing a compartment with alkali to absorb the COz formed and the other without alkali. The difference in the pressure change recorded on the manometers of the two vessels represents the COz absorbed. Fig. 5 shows the O2 consumption when oxalic acid solutions of the pH values recorded are mixed with ferrous sulfate. Fig. 6 shows the accompanying CO, production. It may be seen that at the less acid reactions the oxygen consumption proceeds rapidly until an amount of oxygen approximately equivalent to the ferrous iron present is consumed, and then stops. A glance at Fig. 6 reveals that at these reactions there is no COz production. The produc- tion of CO2 reaches a maximum in the neighborhood of pH 2.9 and then drops off rapidly. It will be seen immediately that the CO2 production occurs over just that range of acidity where mon- ovalent oxalate ion

0

NwcNO

/ \ HO O-

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C. V. Smythe 263

exists. This result agrees splendidly with our considerations. In ferrous oxalate

the iron should be readily oxidized, but there is no reason for the oxalate radical to be oxidized. This is an ordinary &membered ring and such rings are usually quite stable. This situation is not changed even if in excess of oxalate another molecule of oxalate is attracted by the residual valence of the iron. If, however, we have the monovalent oxalate ion present, the compound formed with ferrous ion is

0 0 //

C-OH HO& I

‘.., /’ ._, ,/ C-O-Fe-O-C

This will pass through a series of reactions entirely analogous to those discussed for tartrate. After losing two hydrogen ions and three electrons it will form

/O

0

\ c-o----o-c

I “‘, .I’ ., . . ‘.., ./ . . . ...’ I

C-0-Fe+-O-C \ //

0 0

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264 Mechanism of Iron Catalysis

and there is no difficulty in seeing why this should lose COZ. It should be pointed out that the oxygen uptake in excess of the iron equivalent in this case is very much slower than in the case of tartrate. One sample at pH 2.9 was run for 22 hours. During this time the COZ production was only 184 c.mm. The 02 con- sumption, in excess of the iron equivalent, for the same period was 40 c.mm. It should be noted that oxalic acid produces 4 molecules of CO2 for each molecule of O2 consumed. The small oxygen

consumption (40 c.mm. instead of !?! = 46 cmm.) may be par- 4

tially due to the fact that not all the iron was in the ferric condition at the end.

Since, as we have seen, ferric oxalate is a stable substance, it follows, that, to the extent that this compound is formed in the oxidation of other organic compounds, the iron will be removed from the realm of active catalysis. In such cases the catalyst may be said to deteriorate during its action.

DISCUSSION

The important thing to be grasped from the above pages is that in a consideration of electron structure and residual charges on atoms may lie the explanation for that vague term “activation” of hydrogen. According to this concept the hydrogen separates from the molecule as hydrogen ion and the process of activation is simply a process of increasing the acid dissociation constant of the group involved. This is accomplished by the iron in the manner described.

Earlier in the paper we mentioned the fact that a number of workers consider that the ferrous iron functions: in such reactions as we have described, by forming an iron peroxide which then acts as the oxidant. The present paper throws no direct light upon this problem; however, we may point out that if molecular oxygen accepts the electron which ferrous iron gives up on being oxidized, it would seem quite necessary that there be a transient compound between them during the passage of the electron. This compound would be an iron peroxide. Such a compound would in no way change our scheme of oxidation and its existence is a matter out- side the scope of this paper.

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C. V. Smythe 265

SUMMARY

It is shown that if ferrous iron is obtained in an unionized compound, it is readily oxidized to ferric iron by the oxygen of the air regardless of the acidity of the solution. An explanation in terms of electron structure is advanced.

The catalytic effect of ferrous and ferric iron in certain oxida- tions is discussed. dn interpretation of the mechanism is proposed.

A suggestion is made concerning the meaning of the term activa- tion of hydrogen.

The author is indebted to Dr. L. Michaelis, in whose laboratory this work was carried out, for many helpful suggestions and for constant advice.

BIBLIOGRAPHY

1. Smythe, C. V., and Schmidt, C. L. A., J. Biol. Chem., 88, 241 (1930). 2. Fredenhagen, C., 2. Chem., 29,296 (1902). anorg. 3. Kolthoff, I. M., 2. ZL. allg. Chem., 110, 143 (1920). anorg. 4. Manchot, W., Ber. them. Ges., 46, 2869 (1912). 5. Baudisch, O., and Davidson, D., J. Biol. Chem., 71,501 (192627). 6. Wieland, H., and Franke, W., Ann. Chem., 464, 101 (1928). 7. Wieland, H., and Franke, W., Ann. Chem., 467, 1 (1927). 8. Wieland, H., and Franke, W., Ann. Chem., 476, 1 (1929). 9. Manchot, W., Z. Chem., 27, 420 (1901); 34, 2479 (1901); anorg. Ann.

Chem., 326,93, 105 (1902-03); 460,179 (1927). 10. Warburg, O., Biochem. Z., 162, 479 (1924). 11. Goard, A. K., and Rideal, E. K., Proc. Roy. Sot. London, Series A, 106,

135 (1924). 12. Latimer, W. M., and Porter, C. W., J. Am. Chem. Sot., 62,206 (1930). 13. Kauffmann, H., Die Valenzlehre, Stuttgart (1911). 14. Latimer, W. M., J. Am. Chem. Xoc., 61, 3185 (1929). 15. Fenton, H. J. H., J. Chem. Sot., 66,899 (1894).

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C. V. SmytheCATALYSIS IN CERTAIN OXIDATIONS

THE MECHANISM OF IRON

1931, 90:251-265.J. Biol. Chem. 

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