160 OXIDATION O F BENZENE

THE OXIDATION OF BENZENE BY HYDROGEN PEROXIDE AND IRON SALTS

BY J. H. BAXENDALE AND J. MAGEE Chemistry Department, Manchester University

Received 20th May, 1952

A quantitative determination of the products of the oxidation of benzene by Fez+ -I- H202 in dilute solution has shown that phenol and diphenyl only are formed. The variation in the amounts of these products in different conditions has led to the conclusion that the phenyl radical does not react with hydrogen peroxide, but may be reduced by ferrous ion to benzene and oxidized by ferric ion to phenol. It is possible that all the phenoI is produced by the latter reaction and not by the combination of phenyl and hydroxyl radicals usually assumed.

In the presence of oxygen far more oxidation of ferrous ion and benzene occurs than can be accounted for by the hydrogen peroxide alone. This is explained in terms of reactions of phenyl hydroperoxide and phenyl hydroperoxyl radicals, the latter being produced from phenyl radicals and oxygen.

The original investigations of Fenton on the oxidation of hydroxy compounds by ferrous salts and hydrogen peroxide have been considerably extended in recent years to cover the reactions of this reagent with a wide variety of compounds. The observation that the reagent will initiate vinyl polymerization as well as oxidize certain substrates and that these two reactions can be competitive 1 led to the conclusion that the same entity was responsible for both. The initiation of polymerization is strong evidence for the participation of radicals and it is now generally accepted that the primary reaction involves the production of hydroxyl radicals as follows :

Fez-’- -t H202 - Fe3f + OH- I- OH. (0) k0

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J . H. B A X E N D A L E A N D J. MAGEE 161

The hydroxyl radical thus produced may add to a vinyl double bond and lead to polymerization or abstract an electron or hydrogen atom from a substrate molecule and lead to its oxidation.

The work of Merz and Waters and of Kolthoff and Medalia has shown that the oxidation reaction may take one of three courses. In some cases, e.g. acetate, acetone, C1-, although it can be shown that the substrate is attacked by the hydroxyl radical there is no oxidation. This arises because the radical formed in the step

(s) H2R + *OH + HR* + H2O X- + *OH -+X + OH-

readily reacts with ferrous ion and reverts to the original substrate. HR. + Fez+ + H+ + H2R -1- Fe3+

X + Fez+ + X- -1- Fe3+. Here the substrate acts merely as a radical transfer agent. With other sub-

strates the amount of oxidation is determined entirely by the competition between H2R and Fe2+ for hydroxyl, i.e. by reaction (s) above and the reaction

(1) In this case the fate of the radical HR- is either

Fez+ + *OH + Fe3+ + OH-.

or HR* + HR. + HR-RH HR* + *OH + HROH.

This has been called non-chain oxidation by Merz and Waters,2 who consider that tertiary alcohols, esters and many aromatic compounds react in this way.

Finally there are systems in which the amount of substrate oxidized is greater than the apparent amount of hydroxyl radical reaction in step (s). This indicates that latter reaction initiates an oxidation chain which consumes hydrogen peroxide. Alcohols, hydroxy acids, amino acids are compounds which behave in this way. The catalytic decomposition of hydrogen peroxide by ferrous ions may also be considered in this class. The details of the mechanism of chain oxidation of most substrates are still in doubt. One possibility is that after reactions (0) and (s) above, the radical HR* reacts with hydrogen peroxide

(P‘) and the hydroxyl radical continues the chain. On the basis of reactions (u), (s), (1) and (p’) above it can be shown that the consumption of ferrous ion (AFe2+) and hydrogen peroxide (AH202) are related to the reactant concentrations approxi- mately by the equation

(A) Merz and Waters 2 found this equation to hold for the chain oxidation of several substrates and adduced this as evidence for reaction @’). However, a possible alternative to (p’) is

HO- + Fe3-1- + HR* + HROH + Fez+. Here the ferrous ion is regenerated and continues the chain. Such a reaction has been shown to occur when hydrogen peroxide is the substrate by Barb, Baxendale, George and Hargrave 3 and has been suggested as an alternative to (p’) for organic substrates by these authors and by Kolthoff and Medalia.4 It is to be noted that provided all the radical HR. reacts as in (p) then eqn. (u), (l), (s) and (p) will also lead to eqn. (A), so that measurement of AFe2+ and AH202 will not dis- tinguish between reactions (p’) and (p). Moreover, the inability of added ferric ion to affect AFe2+ and AH202 does not preclude the reduction of Fe3+ by radical HR. as suggested by Men and Waters,s for if all HR. reacts in this way increase of ferric ion concentration will not result in any change.

HR* -1 H202 -+ HROH + *OH

2 AH202/(2AH202 - AFe2+) = 1 + k1[Fe2+]/ks[RH2].

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162 O X I D A T I O N O F B E N Z E N E

There are two ways by which it is possible in principle to distinguish between reactions (p’) and (p). A kinetic analysis shows that the rate of ferrous ion oxida- tion in such systems is unaffected by the addition of ferric ion if (p’) is operative, but will be decreased if reaction (p) occurs. To test this the high velocity of reaction (0) requires the use of very low reactant concentrations (- 10-4 M). However, this introduces practical difficulties, since at these low concentrations in the presence of air with organic substrates there occurs an induced reaction with oxygen which upsets the above kinetic schemes. The details of these induced reactions have been discussed by Barb et al.3 and by Kolthoff and Medalia 4 and the “ dilution effect ” observed by Merz and Waters 2 can be attributed to such reactions. The only substrate free from these objections is hydrogen peroxide itself and here the evidence is in favour of reaction ( p ) .

Another approach is to use conditions where alternative reactions of the radical HR* occur to an appreciable extent in which case a change in the concen- trations of hydrogen peroxide or ferric ion will affect the course of the reaction depending upon whether (p’) or (p) occurs. One such alternative might be, for example,

(4 Merz and Waters in their analysis of the products of chain oxidations did not report the presence of compounds corresponding to HRRH in any of the many substrates examined. It would appear therefore that reaction (d) does not occur to any appreciable extent in the concentration conditions (- 0.02 M Fe2+ and H202) used by them.

HR* + HR* + HRRH.

EXPERIMENTAL

MATERIALS.-PerchlorateS were used throughout. Ferrous perchlorate was pre- pared by dissolving spectroscopic iron in A.R. perchloric acid in an initial atmosphere of oxygen-free nitrogen. Ferric perchlorate was obtained from the ferrous salt by oxidation with pure hydrogen peroxide. The latter was Laporte’s 85 % stabilizer-free product and was found to give results identical with those from a distilled sample of the same material. A.R. grade benzene was subjected to three recrystallizations before use. A.R. phenol was used without further treatment. Diphenyl was purified by recrystallization and sublimation and its purity checked by melting point,

REACTION VESSEL.-Reactions were carried out in a 750 ml round-bottomed flask fitted with two side arms and with a tap for connection to a vacuum line. Initially the flask contained 250 ml of ferrous perchlorate solution of known strength and in one side arm was pipetted 2 ml of benzene. The other side arm consisted of a ground glass joint holding a glass rod which supported a small glass tube containing a known amount of hydrogen peroxide. Deaeration was carried out by pumping the solution hard while vigorously stirring with a glass-covered magnetic stirrer. During the pumping process the benzene was held frozen in the side arm and was itself deaerated by successive melting and freezing. It was found that three pumping periods of 3 min each with intervals of about 5 min were sufficient to give results consistent with the removal of all the oxygen. The test of this was the oxidation balance discussed below. It was established by estim- ation of the hydrogen peroxide and ferrous perchlorate solutions before and after that the loss of these materials during the pumping procedure was negligible.

After deaeration the benzene was run into the aqueous solution and stirred vigorously to give a saturated solution. The oxidation was then begun by dropping the hydrogen peroxide tube into the solution, again with vigorous stirring.

ANALYSES.-Ferrous ion was estimated colorimetrically as the ferrous iris-(dipyridyl) complex ion using the procedure described elsewhere.3 This method, however, is not reliable when the ferric ion concentration exceeds about 5 x 10-3 M. In such cases we have used the reduction of ferric tris-(dipyridyl) by ferrous ion to estimate the latter. A 2 x 10-3 M solution of ferric iris-(dipyridyl) in 0.01 N sulphuric acid was prepared by oxidizing the corresponding ferrous complex ion solution with lead dioxide. One ml of this solution was run into a convenient amount of the ferrous solution to be estimated. The optical density of the ferrous tris-(dipyridyl) solution thus produced was immediately measured. Owing to the fading of the colour in the strong acid solutions used in this

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J . H . B A X E N D A L E A N D J . M A G E E 163

procedure, it was necessary to observe the optical density over a period of lime and extra- polate back to the time of mixing. This extrapolation did not extend to more than 5 % of the total optical density and checks with known ferrous ion concentrations showed that the method was just as accurate as the direct procedure used at low ferric ion concentra- tions. The extent of the fading can be decreased considerably by cooling the solutions and it would seem that if the analogous rhenanthroline complexes are used it might be reduced even further since these are more stable than the dipyridyl ones. We propose to examine this possibility.

When the extent of oxidation of ferrous ion was only a few per cent. it was necessary to measurc the amount of ferric ion produced in order to get the required accuracy. This was done using the usual thiocyanate + acetone method. It was found that in the presence of large amounts of ferrous ion the optical density of the ferric thiocyanate increased slowly with time, due no doubt to oxidation of ferrous ion. This was so slow that back extrapolation was unnecessary if the measurements were made immediately after mixing.

In the analysis of the reaction mixture for phenol and diphenyl the solution was first extracted with 50 ml of spectroscopically pure hexane. It was found by blank experi- ments that this was sufficient to remove all the diphenyl and most of the benzene, whilst leaving all the phenol in the aqueous phase. A second extraction with 20 ml of hexane removed the benzene almost entirely. The phenol in the resulting aqueous solution was then extracted successively with 50 ml and 20 ml of ether, the ethereal extracts being made up to 50 ml in a graduated flask. This solution was then examined spectroscopically and the concentration of phenol calculated from the optical density at 273.5 mp. On our particular spectrophotometer phenol shows a peak at this wavelength with an ab- sorption coefficient e of 2170. Diphenyl in hexane shows maximum absorption at 247 mp with E = 17,400 but the hexane extract obtained as above cannot be measured directly since benzene is also present. Benzene was removed by evaporating it off with the hexane from a known amount of the solution in vacuo. The following procedure was shown by blank experiments to give reproducible results. To a known amount (10-20 ml) of the hexane extract was added 1 ml of a 20 % solution of liquid paraffin in hexane. This mixture was subjected to the usual vacuum degassing procedure using a vacuum system at about 10-3 mm. The volatiles were then allowed to distil into a trap at - 80" C over a period of 40-45 min. The residual liquid paraffin + diphenyl mixture was then taken up in a known amount of hexane and examined spectroscopically.

In both the phenol and diphenyl estimations it occasionally happened that traces of benzene remained and in each case the absorption of benzene must be corrected for. This was done by measuring the optical densities at 255 mp where (with the slit widths used here) benzene shows a maximum. Then with a knowledge of the e of diphenyl and phenol at 255 mp and that of benzene at 247 mp and 273.5 mp a solution of simple simul- taneous equations gives the required concentrations.

These procedures were checked on aqueous solutions containing known amounts of diphenyl and phenol and gave results which showed a maximum scatter of about 5 % in both cases.

DISCUSSION

BENZENE oxImTIoN.--Tt was observed by Merz and Waters 6 that in addition to the phenol and catechol previously detected as products of the oxidation of benzene by Fenton's reagent, diphenyl is also formed. This is attributed to reaction (d). These authors concluded from a quantitative analysis of the reaction that the oxidation goes via a non-chain mechanism, since they found

2AH202/(2AH202 - AFe2+) = 2 + kl[Fe2+]/ks[RH2] (B)

which is consistent with the occurrence of reactions (o), (l), (s) and ( t ) or (d). The kinetics of the reaction scheme involving (It) simultaneously with either ( t ) or ( p ) , which must apply to the benzene oxidation in view of the simultaneous occurrence of phenol and diphenyl, are too complex to yield a solution easily tested. The system is interesting in that an analysis of the oxidation products in different conditions will give information about reactions ( p ) and (p') since there is an alternative reaction (d ) for the radical HR..

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164 O X I D A T I O N OF B E N Z E N E

In the present investigation we have kept the reactant concentrations low (N 10-4 M) so as to avoid the further oxidation of the initial products which is probably the source of the catechol previously reported. Using saturated aqueous solutions of benzene at 25" C with these low concentrations of ferrous ion and hydrogen peroxide, we have not been able to detect any products other than phenol and diphenyl, and the oxidation balance in various conditions (table 1) shows that the other product can only be present in negligible amounts.

OXIDATION BALANCE.-^ the estimation of the oxidation products under various conditions we have observed that the presence of air has a considerable effect on the nature and amount of these products. Thus in the absence of air, phenol and diphenyl are produced, but in air only phenol could be detected. Moreover, in the presence of air the total amount of oxidation is too large to be attributed to hydrogen peroxide alone (expt. 8 and 9, table l), and it must be concluded that some oxidation by oxygen also occurs. The reaction in these conditions is discussed in more detail below.

TABLE 1 . 4 - 1 0 M HClO4, SATURATED BENZENE SOLUTIONS AT 25" C

Amounts in moles/l. x 105

PhOH total oxidation (equiv. /2)

H Z 9 2 (initial) conditions

1 20-5 10.4 2.8 1-04 9.0 9.36 2 20.9 10.2 2.8 1.08 9.0 9.32 3 20.6 9.5 2.7 1-48 9.0 9.24 4 20.1 7.5 1.9 3 4 9.1 9-20 5 20.3 1.5 0-50 7.6 8.9 9.20 6 40.0 10.8 2.8 0.88 9.1 9.32 7 122.0 11.7 2.5 0.80 9.2 9-32

evacuated 1 20.4 0 6.6 16-8 9.32}in air 8 20.4

9 123.0 47.6 0 10.9 34.7 13.3 -

2.1 1.04 9;:p} evacuated 10 8.3 8.3 11 8.3 8.3 2.4 0.96 -

In table 1 typical examples of oxidation balances are given. It will be seen that in the absence of air the total amount of the products ferric ion, phenol and diphenyl are equivalent to the initial hydrogen peroxide to within 5 % or less, which is about the error in the determinations. We can conclude therefore that to this accuracy other oxidation products can be neglected. This equivalence is useful in that it also serves as a check on the absence of oxygen from the system. We will consider fist the observations made in the absence of air.

that the existence of two products phenol and diphenyl makes it possible to dis- tinguish between reactions (p') and (p). Thus, if (p') occurs it is to be expected that increase in the hydrogen peroxide concentration will favour the production of phenol at the expense of diphenyl. We have increased the peroxide concentra- tion ten times (expt. 10, 11, table l), and this has very little effect on the relative amounts of the products. It would appear then that reaction (p') does not occur in the present system.

EFFECT OF FERRIC ION CONCENTRATION.~n the other hand, the data in table 2 show that the ferric ion concentration has a considerable effect in decreasing the amount of diphenyl relative to phenol. Thus the presence initially of ferric ion at the same concentration as the ferrous ion increases the amount of phenol from 28 % to 36 % (expt. 12, 13). The amount of phenol increases with the ferric ion and it will also be seen that the amount of ferrous ion oxidized simultaneously decreases. At the high ferric ion concentration of expt. 15 the oxidation of ferrous ion only constitutes about 8 % of the total oxidation so that the reaction is almost

EFFECT OF HYDROGEN PEROXIDE CONCENTRATION.--It was pointed Out above

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J . H . B A X E N D A L E A N D J . M A G E E 165

TABLE 2.-sATURATED BENZENE SOLUTIONS AT 25" c, IN THE ABSENCE OF AIR

Initial [Fez+] = 20.4 x 10-5 M ; initial [H202] = 9-20 x 10-5 M ; amounts in moles/l. x 105

12 0.0 10.2 2.8 1.08 0.1 N acid 13 10.4 9.5 2.6 1.48 0 1 N acid 14 86.5 7.52 1.9 3.40 0-1 N acid 15 864-0 1-48 0.50 7.6 0.1 N acid 16 0 10.0 2.5 1.28 3 x lO-4N acid 17 0 18.2 0 0 0.02 N alkali 18 0 9.9 3.3 0.76 0.04 M Na2HP04

expt. Fd+ added Fez+ oxidized Ph2 PhOH conditions

a pure catalytic oxidation of benzene by hydrogen peroxide. This is clearly a chain oxidation process.

It will be noted from the data in table 2 that for expt. 14 and 15 the amount of ferrous ion oxidized is less than the molar amount of hydrogen peroxide present. Now with reaction (0) as the primary reaction, this can only result if ferrous ion is in some way regenerated, or hydrogen peroxide is consumed by another reaction. The reaction of ferric ion itself with hydrogen peroxide is too slow at these acidities to account for any removal of hydrogen peroxide and we have eliminated the possibility of the reaction of hydrogen peroxide with radical intermediates in step (p'). It seems probable then that the ferrous ion consumed in step (0) is reformed in another reaction. Moreover, this process becomes more important as the ferric ion concentration increases and simultaneously the phenol production increases relative to diphenyl. All these observations are good evidence for the fact that ferric ion oxidizes the phenyl radical to produce phenol and ferrous ion, i.e. that reaction (p) above is important in the present system. This oxidation may occur in several ways involving either the hydroxyl ion, free or in an ion pair :

Fe3f + OH- + Ph- -+ PhOH + Fez+ (a) FeOH2-I- + Ph- -+ PhOH + Fe2+ (6)

(4 (4

or involving a water molecule H 2 0 Fe3+ + Ph- -+ Ph+ + Fe2+ - PhOH 4- H+ + Fez-'-

or Fe3+. H20 -1- Ph- + PhOH + H-1- -1- It would be difficult to distinguish between (c) and (d), but in principle it is possible to determine whether (a) and/or (b) are operative by a study of the effect of acidity on the phenol production, making use of the known equilibrium constant for the reaction

Fe3-l- -1- OH- -2 FeOH2+. We have not examined this aspect of the reaction in any detail, but expt. 16 in table 2 was done in 3 x 10-4 N acid compared with the usual 0.10 N. At this acidity about 95 % of the ferric ion is present as FeOH2-C. compared with about 5 % at 0.10 N. It will be seen that there is little effect on the relative amount of phenol produced although the concentrations of Fe3+ and FeOH2+ have changed about twentyfold and that of HO- by a factor of 3 x 103. This must mean either, that reaction (a) does not occur and the rate constants for (6) and (c ) and/or (d) are about equal, or, that (c) and/or (d ) occur but the rate of formation of the ion pair FeOH2+ is slow compared with these reactions. Expt. 17 of table 2 was done in N/50 alkali to find the effect of higher HO- concentrations. It will be seen that. here there is no benzene oxidation at all. This can probably be attri- buted to the hydroxylated ferrous ion which is probably present, reacting more quickly with the hydroxyl radical than does the free ferrous ion, so that the reaction with benzene is prevented.

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166 O X I D A T I O N O F BENZENE

The production of phenol by ferric ion oxidation of the phenyl radical raises the question as to whether any phenol is formed by the reaction

Ph* + *OH --t PhOH

as has been assumed by Merz and Waters6 and by Stein and Weiss.7 The data of expt. 15 can be used to show that very little is formed in this way in the presence of a large amount of ferric ion. Thus the formation of a molecule of phenol or diphenyl by the radical combination reaction must be accompanied by the oxida- tion of at least two ferrous ions in the production of two hydroxyl radicals. Hence, allowing for the diphenyl in expt. 15 only 3 % of the phenol can be formed by the radical combination reaction, and it may be less if any ferrous ion is oxidized by hydroxyl radical (reaction (I)) or other means (see below). It is possible that although no ferric ion is present initially in expt. 12, the ferric ion produced is responsible for the whole of the phenol. One way to check this would be to remove the ferric ion as it is formed. We have attempted this in expt. 18 by having phosphate present to complex the ferric ion. However, it can be seen that an appreciable amount of phenol is still formed under these conditions which still leaves the question open.

It is clear from these observations on the part played by ferric ion in this system, that care must be exercised in drawing conclusions from a comparison of the action of Fenton’s reagent on a substrate with that of other sources of hydroxyl radicals, e.g. ionizing radiations.

table 2 is that as the ferric ion is increased the apparent amount of phenyl radicals produced (2 x Ph2 + PhOH) also increases. It is difficult to see how ferric ion can interferc with reactions (o), (1) and (s) which determine the amount of phenyl, and it seems more probable that the amount produced is the same but that there is another reaction of phenyl radicals which does not lead to diphenyl or phenol and which competes with the ferric ion + phenyl radical reaction. This reaction is probably

(r)

Support for this is given by expt. 12 and 15. Thus if we assume that all the phenol in expt. 15 arises from the oxidation of phenyl by ferric ion and hence produces an equivalent amount of ferrous ion, the total amount of ferrous ion oxidized in reactions (0) and (1 ) is 9.1 x 10-5 moIes/l. Now this is about the same as the amount of hydrogen peroxide present initially which will react with ferrous ion in reaction (0). Hence very little ferrous ion is consumed by reaction ( 1 ) in these conditions. Nor is any more likely to react in this way in expt. 12, since here the mean ferrous ion concentration is smaller than in expt. 15 which favours reaction (s). However, in expt. 12 more ferrous ion is oxidized than can be accounted for by reaction (0) alone, even if it is assumed that none is regenerated in the production of phenol. Hence another reaction must be postulated which oxidizes ferrous ion, and this i s satisfied by reaction (r) above.

It would have been additional evidence for the absence of reaction ( l ) , and hence support for (r), if it could be shown that a variation in the benzene concen- tration had no effect on the reaction. This is technically difficult in vacuo, and as an alternative approach we have examined the effect of ferrous ion concentration. The data are given in table 3. Here we have listed the amount of ferrous ion oxidized in excess of that oxidized in reaction (0) (i.e. total ferrous ion oxidized- initial peroxide). We have seen, however, that some ferrous ion is regenerated by reaction ( p ) in the production of phenoi, so that to get the total amount of ferrous ion oxidation we must allow for this. We do not know exactly how much of the phenol arises from this reaction and how much from the combination of radicals, but we can take the two extremes and assume first that no ferrous ion is regenerated (a in table 3), and secondly that it is formed simultaneously with

EFFECT OF FERROUS ION CONCENTRATION.-Another feature Of eXpt. 12-1 5 in

Fe*+ + Hf + Ph* -+ PhH + Fe3+

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J . H . B A X E N D A L E A N D J . MAGEE 167

Now the variation in this ferrous ion every phenol molecule (b in table 3). oxidized may arise from the competition between

Fez+ + *OH + Fe3f + OH- (1) and PhH -t *OH + Ph* + H2O (4 The amounts (a) and (b) of table 3, which we call x, would then be the ferrous ion oxidized in reaction (1). Denoting the initial hydrogen peroxide concentration by a, then on this view (a - x) is the amount of *OH reacting in (s). Hence we should find approximately that

x/(a - x) = kl[Fe2-'-]/ks[PhH]. Since [PhH] is constant, x/ (a - x) should then be proportional to the mean ferrous ion concentration. However, the data in (a) and (b) of table 3 show that x/(a - x ) varies only fourfold and threefold respectively for a change of about twentyfold in the ferrous ion concentration. It is clear therefore that the above competition reactions cannot account for the observations. This would be consistent with the previous conclusion that reaction (1) does not occur appreciably in these conditions.

TABLE 3 . 4 - 1 0 N HCIO4, SATURATED BENZENE SOLUTIONS AT 25" C IN THE

Amounts in moles/l. x 105 ABSENCE OF AIR

Fez+ oxidized in excess of reaction (0)

expt initial Fe2k initialjH202 PhOH final Fe3+ (a) ( b )

19 20.48 9-36 1-04 10.4 1 -04 2.08 20 40.0 9.32 0-88 10.8 1 -48 2.36 21 122.0 9-32 0.8 1 11-7 2.38 3-19 22 41 5.0 9.28 0.70 12.2 2.92 3-62

A feature of the data in table 3 is that the large variation in ferrous ion concen- tration has a comparatively small effect on the products. If all the phenol arises from the oxidation of phenyl radicals by ferric ion, it might be expected that since this is in competition with reaction (r) above, an increase of twentyfold in the ferrous/ferric ion ratio would lead to a considerable decrease in the phenol. The fact that this is not the case is some indication that phenol is formed, at least partially by the radical combination reaction. It is unfortunate that the com- plexity of the system makes it impossible with the present data to carry out a quantitative analysis on the basis of the proposed reactions.

organic product is phenol. Diphenyl, if present at all, must be there to less than 10-8 mole/]. The amounts of products in various conditions are listed in table 4.

REACTIONS IN THE PRESENCE OF OXYGEN.-h air or oxygen the Only detectable

expt

23

24 25 26 27 28 29 30 31 32

TABLE 4.-

initial H2Oz

9.36

9.32 2.80 1-86 1.86

13.3 13.3 13.3 13.3 13.3

-0.10 M HC104 SATURATED BENZENE SOLUTIONS AT 25" C amounts in moles/l. x 10s

initial Fe

20.5

20.4 20.3 4.10 2.10

123.0 123.0 125.0 124.0 209.0

initial Fez+ Fe3 + oxidized PhOH conditions

0

0 0 0 0 0

910 0

910 0

10.4 1-04 evacuated

20.4 6.60 air 19-2 4.60 air 4.10 1 -46 air 2-10 0-84 air

47.5 10.9 air 40.0 9.4 air 89.0 15-5 oxygen 59.5 11.3 oxygen 83-0 11.6 oxygen

(-I- 2.8 Ph2)

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168 O X I D A T I O N OF BENZENE

It will be seen that always the total amount of oxidation exceeds the amount of hydrogen peroxide present initially. This can be interpreted in terms of reactions previously discus~ed.3~ 4 Thus phenyl radicals produced by reactions (0) and (s) above react with oxygen as follows :

and the lack of diphenyl can be attributed to this reaction, being much faster than the radical dimerization. Subsequent reactions of the hydroperoxide radical are probably

Ph* + 0 2 + PhOp

H+ + PhOz* + Fe2+ -+ Ph02H + Fe3+ (4

H+ + PhO* + Fe2+ -+ PhOH + Fe3+ The increase in oxidation in going from air to oxygen (expt. 28, 30) indicates that the phenyl radical is concerned in other reactions, and since increase of ferrous and ferric ions both decrease the extent of oxidation (expt. 30, 31, 32), reactions

PhOzH + Fe2+ + PhO. + Fe3+ + OH-

Fe2+ + Ph- -t Hf -+ Fe3+ + PhH Fe3-+ + Ph- + OH- + PhOH + Fez+

which have been considered previously, will account for the observations. It can be seen on the basis of these reactions each hydrogen peroxide molecule

reacting in (0) cannot yield more than one molecule of phenol. However, expt. 25 and 30 show that this is not always the case. It follows that other chain pro- pagation reactions must be present. One possibility is that instead of (h) above we have

Ph02H + Fez+ -+ PhO- + Fe3+ + *OH so that the hydroxyl continues the chain through reaction (s) without further con- sumption of hydrogen peroxide. Another alternative is that there are present the chain reaction steps which are known to be important in the oxidation of olefins,8 viz. :

PhO* + PhH -+ PhOH + Ph* PhOy + PhH +- Ph02H + Ph*

A further possibility is the hydrolysis of phenyl hydroperoxide to regenerate hydrogen peroxide

PhO2H 3- H2O -+ PhOH + H202.

In certain of the experiments in table 4 where the amount of phenol produced is high, it will be seen that as many as seven ferrous ions are oxidized for each phenol molecule produced. However, from the above reactions four is the maximum possible. The origin of this discrepancy is not clear, but we suggest two possi- bilities. First that phenol is being attacked by the hydroxyl radical in reaction (s) and is oxidized by similar reactions to those above through o-benzohydroquinone, o-benzoquinone and ultimately to muconic acid ; or that the radical PhO-, which may be considered as

o = = or o=D- /--/

may take up oxygen and proceed similarly. W e have not been able to detect the presence of quinones or hydroquinones in the final products, but muconic acid, if present, would not be detectable by our method of analysis.

THE FERRYL IoN.-It has been pointed out 39 4 that the reactions of Fenton’s reagent can be equally well attributed9 to the ferryl ion FeOH3+ or FeO2+ as to the hydroxyl radical, and this is true in the present system. The ferryl ion, formed either directly in reaction (0) or in equilibrium with NO* as follows

Fe3+ + HO -2 FeOH3+--f FeO2+ + H+

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J. H. B A X E N D A L E A N D 3 . MAGEE 169

can be written for HO- in all the reactions considered above, and most of the observations could also be interpreted in terms of ferryl ion reactions. Thus the increase in the amount of phenol produced by added ferric ion could be explained by an increase in ferryl ion concentration through the above equilibria, together with the assumption that phenol can be formed by

FeO2+ + PhH -+ PhOH + Fez+.

The enhanced attack of benzene in these conditions, previously explained in terms of the competition between reactions ( r ) and (p), can be accounted for in the same way. Moreover, the inability of phosphate or low acidites to affect the amount of phenol might be expected if ferryl ion and not ferric ion was the reactive species. However, such interpretations are not specific evidence for the ferryl ion in view of the alternatives which have been given above in terms of the hydroxyl radical. We prefer this more usual formulation, although it may need to be amended should more specific evidence for the existence of the ferryl ion be forthcoming.

RADICAL ADDITION REAcnoNs.-Although in the above schemes we have assumed that the hydroxyl abstract hydrogen from benzene to give a phenyl radical which may dimerize or be oxidized to phenol, there seems to be good evidence that radicals tend to add to benzene and substituted benzenes10 In the present system there are two such possibilities. First, instead of abstracting hydrogen the hydroxyl radical may add to benzene to give the radical I. If this occurs then phenol can be produced by oxidation of I either by HO- or by Fe3+.

If I is the only product one must then assume diphenyl to be formed by dimeriza- tion to I1 and subsequent dehydration to diphenyl. The second possibility is that hydrogen abstraction may occur initially to give a phenyl radical, but this may add to benzene to yield 111, or yield phenol by the reactions discussed pre- viously. Further oxidation of I11 to diphenyl by the hydroxyl radical may then follow. Assuming all these reactions were possible, we see nothing in our experi- mental data to support or reject these radical addition reactions. A detailed kinetic analysis would be required before any conclusions could be drawn. Such an analysis is difficult with the hydrogen peroxide -+ iron system, but we hope to clarify the position by studying the oxidation of benzene induced by the photolysis of hydrogen peroxide.

1 Baxendale, Evans and Park, Trans. Faraday Soc., 1946,42, 155. 2 Merz and Waters, J. Chem. SOC., 1949, S15. 3 Barb, Baxendale, George and Hargrave, Trans. Faraday Soc., 1951, 47,462. 4 Kolthoff and Medalia, J. Amer. Chem. SOC., 1949, 71, 3777, 3784. 5 Merz and Waters, ref. (2), p. 519. 6 Merz and Waters, J. Chem. SOC., 1949, 2427. 7 Stein and Weiss, J. Chem. Sac., 1949, 3245. * Bolland, Quart. Rev., 1949, 3, 1. 9 Bray and Gorin, J. Amer. Chem. SOC., 1932, 54, 2124.

10 Angood, Hey, Nechvabel, Robinson and Williams, Research, 1951, 4, 386.

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