HYDROGEN PEROXIDE + WATER MIXTURES PART 4.-CATALYTIC DECOMPOSIIP'ION OF HYDROGEN PEROXIDE

BY PETER JONES, R. KITCHING, M. L. TOBE AND W. F. K. WYNNE-JONES Physical Chemistry Laboratories, King's College, University of Durham,

Newcastle-upon-Tyne, 1.

Received 21st March, 1958

The catalytic effect of iron perchlorate on the decomposition of hydrogen peroxide in aqueous solution has been studied at mole fractions of hydrogen peroxide up to 0.95 and at temperatures of lo", 25" and 40°C. The reaction is homogeneous, provided that the solution is kept sufficiently acid, with a rate which is found to follow the expression,

- dX,/dt = ax, + bX$, where A', is the mole fraction of iron perchlorate, X, is the mole fraction of hydrogen peroxide and a and b are parameters dependent upon pH, temperature and solvent com- position. In concentrated solutions both a and b are linear functions of X,/(l - A'*,). This is interpreted as arising from the replacement of molecules of water in the solvation shell of the ferric ion by molecules of peroxide to give ferric-peroxy complexes of the type [Fe(H20)5 (H202) ]3+. The kinetics of the reaction are compared and are found to be consistent with the radical mechanism originally proposed by Haber and Weiss for the reaction in dilute solution.

The catalytic decomposition of hydrogen peroxide by iron salts has been the subject of many investigations. The extensive literature has been reviewed briefly by Schumb et aZ.1 and in more detail by Baxendale 2 and Weiss.3 The reaction has frequently been termed a ferric-ion-catalyzed reaction, in order to distinguish it from the reaction in which the ferrous ion is oxidized to ferric by hydrogen peroxide without oxygen evolution.

Interest arose in the mechanism of this reaction after Haber and Willstater had suggested that the free radicals OH and H02 might play an important role in the reactions of hydrogen peroxide and Haber and Weiss 4 discussed this proposal with particular reference to the iron salt + hydrogen peroxide system. Although many important features of the reaction have been elucidated some aspects of the mechanism remain in dispute.

All previous kinetic studies of the reaction have been confined to very dilute solutions of hydrogen peroxide so that the aqueous solvent has been inappreciably altered. In the present work we have studied the reaction under homogeneous conditions, that is with the catalyst present as hydrated or hydroperoxidated ions and not, for example, as colloidal particles, in I3202 + H20 mixtures ranging in composition from dilute aqueous solutions to almost anhydrous hydrogen peroxide.

By comparing the predictions of the theories of the reaction, which have been developed on the basis of studies in dilute aqueous solutions, with our results, we have attempted to find out how far these theories can explain directly, or can reasonably be extrapolated to explain the behaviour in more concentrated solutions. In this connection it is necessary to note that, in this investigation, we have to compare rates of reaction in different solvents and any determination of mechanism from kinetic form must be preceded by a distinction between kin& effccts and medium effects. Analysis of our results suggests that, in this case, the mzdium efkcts are unusually simple and we consider that this arises from the structural similarity of the solvents and the isodielectric property of H202 + H20 mixtures.

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80 DECOMPOSITION O F HYDROGEN P E R O X I D E

Studies of the catalytic decomposition of concentrated hydrogen peroxide solutions are of interest because of the increased practical importance of this material in recent years. The rate of decomposition of commercial High Test Peroxide (about 85 % w/w H202) is very low under normal conditions of storage, but increases very markedly if the solution becomes contaminated with traces of catalytic impurities such as iron salts. A number of studies of the reaction from the practical stability viewpoint are in the technical literature and have been reviewed by Schumb et aZ.1

It has been found that the rate of decomposition of hydrogen peroxide may be substantially reduced, that is, the stability of the material may be improved, by the addition of suitable stabilizers. One aspect of the present study of thecatalytic decomposition of hydrogen peroxide has been to obtain a quantitative basis for assessment of the performance and for determination of the mode of action of various stabilizers. The results of our investigations of stabilizer action will be described in a separate paper.

EXPERIMENTAL MATERIALS

HYDROGEN PEROXIDE.-High Test Peroxide, an 85 % w/w unstabilized aqueous solution, supplied by Laporte Chemicals Ltd., was used as the source of hydrogen peroxide. It was found that, in the present catalytic studies, there was no difference in the behaviour of this material when undistilled, distilled once or distilled three times. Undistilled peroxide was therefore used in all subsequent work except where very concentrated material was required. The method of distillation has been described earlier.

DISTILLED wATER.-Laboratory-distilled water was distilled first from dilute alkaline KMn04 and then from dilute phosphoric acid and stored in Pyrex containers. In the early stages of the work a further two-stage distillation was carried out in a silica still but no difference was observed between the behaviour of three-times and five-times distilled water in catalytic experiments and three-times distilled water was used in all subsequent work.

PERCHLORIC ACID.-A.R. 72 % HC104 was diluted to the required concentration with distilled water and standardized with standard NaOH solution. No difference was observed between the behaviour of undistilled and distilled acid.

FERROUS PERCHLORATE.-This was at first prepared by dissolving pure iron wire in perchloric acid but the reaction is slow in the cold and in hot solution side reactions occur. This method was, therefore, discarded in favour of an electrolytic method developed in collaboration with Dr. M. Fleischmann in these laboratories. The electrolyte was 120 ml of 3N HC104. The anode consisted of a rod of spectroscopically pure iron, supplied by Johnson Matthey and Co., suspended in solution by means of a gold hook which was sealed through the glass to the contact. The cathode was a gold plate located at the bottom of the cell so that the stream of hydrogen produced stirred the liquid. The hydrogen was bubbled through a water trap to prevent atmospheric oxygen diffusing back into the cell. A constant current was passed for the time required to give a solution that was 1 M with respect to both Fe(C104)~ and HC104. The concentration of ferrous perchlorate was estimated both titrimetrically using KMn04, after treatment with SnC12 and HgC12 to ensure complete reduction to the ferrous state, and from the weight loss of the anode. The two methods gave results in very good agreement. After the completion of the electrolysis the solution was centrifuged to remove traces of suspended carbon which presumably were present in the iron. The ferrous perchlorate solution contained an equimolar amount of HC104 to ensure that the oxidation 2Fe2+ -t- H202 --f 2Fe3+ + 20H- did not alter the hydrogen ion concentration of subsequent reaction solutions. More dilute catalyst solutions were prepared from this stock solution by volumetric dilution with distilled water.

APPARATUS

Reaction vessels consisted of Pyrex tubes with a water-jacketed condenser fused to the top and were calibrated to hold 20.01111 of solution. The decomposition of hydrogen peroxide was followed by measuring the oxygen evolved, using gas burettes with water as the containing liquid. It was often necessary to follow a reaction over a considerable change in hydrogen peroxide concentration and in such cases the total amount of oxygen evolved at any point in the experiment was used as a measure of the extent of decomposition.

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P . JONES, R . KITCHING, M . L. TOBE AND W . F. K . WYNNE-JONES 81

In the intervals between rate measurements the evolved oxygen was led into a wash bottle and the water displaced was collected in a measuring cylinder. It was found that the results were unaffected if the reaction vessel was heated to accelerate the decomposition between rate measurements or cooled to slow the reaction and thus to arrive at the required hydrogen peroxide concentration at a convenient time. Volumes of oxygen were measured at room temperature and atmospheric pressure and were considered to be saturated with water vapour ; corrections to s.t.p. were made accordingly. It was found at an early stage that some of the reactions under inves:igation were light-sensitive and nigrosine black dye was added to the water thermostats in all subsequent work.

The precautions necessary in cleaning glass-ware for use with concentrated hydrogen peroxide solutions have been described elsewhere. It was found that, after cleaning in this way initially, reaction vessels could be merely rinsed and then soaked in distilled water between experiments without adversely affecting the subsequent results for a considerable time.

UNITS OF CONCENTRATION

The decomposition of conceiitrated aqueous solutions of hydrogen peroxide involves as reactant one component of a solvent mixture and the mixture is progressively diluted by the reaction product which constitutes the second component of the mixed solvent. The weight and density of the solution are also changing. If we can ignore effects of the changing medium upon activity coefficients and the reaction mechanism, then we can consider the reaction rate in terms of changes in the concentration of species. A problem of units of concentration then arises. As the hydrogen peroxide in a solution containing a particular amount of catalyst decomposes, the mole fraction of catalyst remains constant but its molarity and molality increase. If we wish to examine the variation of rate with change in hydrogen peroxide concentration, at constant catalyst concentration, we must decide which, if any, of the units of concentration is kinetically significant.

Most studies of reactions in solution have been made with the reactants in dilute solution in a solvent which does not participate in the reaction equation and which remains effectively unaltered by the small changes in reactant concentration. It is customary in such cases to express concentrations in moles (litre solution)-1. For ionic reactions it is normally neces- sary to keep activity coefficients constant by working at a constant low ionic strength in order to obtain a strict concentration dependence. For more concentrated solutions it is sometimes possible to make an approximate allowance for the changes in activity coeffi- cients, for example, in correlations of the catalytic activity of acids with the acidity function suggested by Hammett. This approach is limited in application. Although little experi- mental work relevant to the present point has been done it appears that in some cases a particular concentration unit is more suitable, in that the rate constants obtained are independent of concentration on one scale of concentration but not on others. In an investigation of the catalyzed mutarotation of glucose in concentrated aqueous solutions of glucose, Kilde and Wynne-Jones 5 found that a constant value for the catalytic constant of the hydroxyl ion was obtained only when they expressed the catalyst concentration in moles kg-1 of water, i.e. effectively the mole fraction.

In view of these considerations we have decided to express concentrations in mole fractions (m.f.) but sufficient information has been given to enable conversions to other units of concentration to be made.

RANGE OF EXPERIMENTAL CONDITIONS

Previous work on this system has been in aqueous solutions, sufficiently dilute for the supposition that the catalyst remains in homogeneous solution to be established on the basis of existing solubility data. In more concentrated hydrogen peroxide solutions it is necessary to show that, under a particular set of experimental conditions, a homogeneous catalytic reaction is obtained, that is, the catalyst is present as hydrated or hydroperoxidated ions and not, for example, in the form of colloidal particles. There is considerable evidence that, in the range of experimental conditions in which we have worked, a homogeneous catalytic reaction occurs.

At pH = 2 and 25"C, in aqueous solution the solubility of ferric iron is of the order 2 X 10-2m.f. In hydrogen peroxide solutions we would expect a smaller degree of ferric ion hydrolysis but the position is undoubtedly complicated by the formation of peroxidated complexes whose chemical behaviour is not established. Banfield and Hilden investigated the light scattering of solutions of ferric ammonium sulphate in 85 % W/W hydrogen peroxide and by plotting intensity of scattered light against pH for a fked salt concentration they

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a2 D E C O M P O S I T I O N OF H Y D R O G E N P E R O X I D E

obtained a straight line whose intercept gave the pH at which that particular amount of iron was just soluble. Thus they were able to construct solubility against pH curves for iron salts in 85 % hydrogen peroxide. An extrapolation of their values gives a solubility of 2 X m.f. of ferric iron at pH = 2, 25°C. This value is in excess of the maximum concentration that we have used in this work.

Banfield and Milden 6 also investigated the variation of the rate of catalytic decomposi- tion of H.T.P. on the pH at constant iron salt concentration. They found that the rate of decomposition passed through a maximum at about pH 3-3, the pH being measured on the scale recommended by Schumb et aL,7 using the E" for the glass electrode in water and making the pH measurement in a solution diluted tenfold with water. On the pH scale that we have used8 their value would be approximately 2.3. We have confirmed this result using iron perchlorate (not at constant ionic strength) and the form of our results is shown in fig. 1. The rate maximum occurs at pH = 2.4. This observation, coupled with

f I 1 I D

I 2 3

PH FIG. 1.-Graph of - (dXp/dt) against pH for the conditions Xp = 0.71. XF = 1.86 x 10-6,

temp. 25°C.

the solubility data leads us to conclude that, below pH = 2.4, the iron salt, in the concentra- tions that we have employed, is completely dissolved and that we have observed a homo- geneously catalyzed decomposition of hydrogen peroxide. The consistency of the rate data supports this conclusion. We have further shown that, under our experimental conditions, the rate of decomposition is independent of the surface, volume ratio of the containing vessel.

RESULTS Two methods have been used to examine the variation of the rate of decomposition

with experimental conditions. In one method the initial rates of decomposition of samples of the required compositions were determined. In the other method the complete decom- positions of samples of concentrated hydrogen peroxide containing catalyst were followed.

INITIAL RATE MEASUREMENTS

The dependence of the rate of decomposition upon the concentration X, of catalyst has been studied at five mole fractions of hydrogen peroxide (Xp = 0.29, 0.50, 0.66, 0.71, 0.78), at 10, 25 and 40°C and pH = 2.

Two methods of making up reaction solutions were used. In one the separate com- ponents were brought to the reaction temperature, mixed in a standard flask and samples measured out into reaction vessels.

For many reactions it was more convenient to make up the hydrogen peroxide solution in the reaction vessel and to start the reaction by adding catalyst. For very fast reactions the reaction vessel was cooled before addition of catalyst. This was necessary to avoid " fuming off" of the hydrogen peroxide which occasionally occurred due to high local

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P. JONES, R . KITCHING, M. L. TOBE A N D w. F . K . W Y N N E - J O N E S 83 concentrations of catalyst during the addition at higher temperatures. After addition of the catalyst the system was immediately sealed and re-immersedin the thermostat. Measure- ments of the extent of decomposition were commenced immediately and rate measurements could then be made at known values of the hydrogen peroxide concentration after temperature equilibrium had been attained. Initial concentrations of hydrogen peroxide were determined in blank experiments by measuring either the density or refractive index of the solution at 25°C.

Graphs of log (dX,/dt) against log X,, where X, is the mole fraction of H202 and XF is the mole fraction of catalyst, showed that the data fitted no simple reaction order. When - (dX,/dt)/XF was plotted against XF, linear graphs were obtained with positive intercepts on the rate axis showing that the results followed expressions of the form :

- (dX,/dt) = ~ X F + bX$, (1) where a and b are constants at constant X,, pH and temperature.

At suficiently low values of XF, graphs of rate against XF were essentially linear so that extrapolation to XF = 0 yielded values for the rate of decomposition of “ pure ” hydrogen peroxide which agreed with the rates of decomposition obtained in blank experiments in

FIG. 2.-Graphs of - (dXp/dt) against (i) X,, (ii) X’/X, ; XF = 4-05 X 10-6, pH = 2 ; temp. 25OC.

which the solution contained no deliberately added catalyst. Experimental reaction rates have been corrected by subtracting the appropriate value for the rate of “ pure ” decom- position. At 10°C and 25°C the correction is negligible. The ‘‘ pure ” rate probably arises from catalytically active trace impurities in the solution and container surface.

Since concentrated hydrogen peroxide solutions readily supersaturate with respect to oxygen and a 20 ml sample of 85 % w/w H202 may contain 10-20 ml oxygen, an error may be introduced into measurements of oxygen evolution. For moderately fast reactions a steady, reproducible rate of oxygen evolution is obtained but for slow reactions the reproducibility and steadiness are much poorer and the distinction between the kinetic form of eqn. (1) and equation of the type - (dXp/dt) = klXp + k2X~3/2 becomes difficult.

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84 D E C O M P O S I T I O N O F H Y D R O G E N P E R O X I D E

To eliminate this error we have used the " banging method " of continuously relieving supersaturation devised by Bell.9

At pH = 2,25"C, Xp = 0.7 we have examined the effect of changes in ionic strength on the reaction rate by adding various amounts of sodium perchlorate to the reaction solution. A small decrease in reaction rate was observed as the concentration of sodium perchlorate was increased, a change in the ionic strength of the solution from 0.01 to 0-045 decreasing the rate by about 10 %. It seems likely in view of the work to be described in a later paper on the effect of added salts on the rate of decomposition that this effect may be accounted for by the formation of a catalytically inactive FeC10$+ complex.

COMPLETE DECOMPOSITIONS

In these experiments the rate of decomposition was measured at forty or more values of X, in order to define the form of the curve precisely. In fig. 2 the lower curve shows the rate of decomposition, under the conditions X, = 4.05 x 10-6, pH = 2, T = 25"C, plotted

FIG. 3.-Graphs for the determination of the constants al, a2, bl and b2 of eqn. (3).

against X,. No simple reaction order in X, could be obtained to satisfy these results. In the upper curve in fig. 2 the rate data for the same experiment are compared with an expression of the form,

(2) where a and /3 are constants at constant X,, pH and T and X, = 1 - X, is the mole fraction of water, by plotting - (dX,/dt) against X,/X,. The results at high values of A', follow this relation very closely but fall off at lower concentrations. The linear section extrapolates to a positive intercept on the rate axis. This form of dependence was obtained at various catalyst concentrations for the conditions pH = 2, T = 25°C and the values of a and fl obtained. The total amount of oxygen evolved agreed with the initial concentration of hydrogen peroxide to 1 %.

- (dX,/dt) = a + B (Xp/Xw),

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P. JONES, R. KITCHING, M . L. TOBE AND w. F . K . WYNNE-JONES 85

From the variation of rate with XF and Xp separately we may write for the rates of

(3)

where al, a2, bl and 62 are constants at constant pH and temperature and are obtained from the constants a, b, a and /3 of equations (1) and (2) :

decomposition (at constant pH and T) of concentrated solutions of H202,

- (dXp/dt) = (a1 + a2(Xp/Xw) >XF 4- (bi 4- b2(Xp/Xw) ) X F ~ ~

a = (a1 4- a2(Xp/Xw) 1 ; b = ( h + b2(Xp/Xw) 1; a = XF(a1 + hXF); f i = XF(a2 f b2XF).

We may thus obtain values of al, a2, bl and 62 and fig. 3 shows graphs for the determina- tion of these quantities at pH = 2, T = 25°C. Table 1 gives the values obtained under all experimental conditions by fitting the data to the expressions given above using the method

TABLE VALUES FOR THE CONSTANTS IN EQN. (3) AT pH = 2.0, T = 25°C. CONCENTRA- TIONS EXPRESSED IN MOLE FRACTIONS, TIME IN MIN

10°C 25°C 40°C

complete initial rate decompositions

3.79 & 0.30 22.0 f 0.8 23.1 f 1.0 71.7 f 2.9 uz 1.58 f 0.13 6.77 f 0.69 9.2 f 0.5 56.5 f 1.6 b~ (2.58 f 0.11) x 105 (3.53 f 0.14) X 106 (2.92 -+ 0.17) X 106 (3-20 f 0.09) x lo7

of least squares. In fig. 3, a is plotted against XF rather than (./IF) against X', the straight line through the origin obtained implying bl = 0. At 10" and 25°C values of bl are zero within the experimental error. At 40°C values of bl - b2 are obtained but this may be because the rate of " pure " decomposition is now significant and that catalyst promotion effects become significant. If we suppose that bl = 0, which is true within the experimental error at 10 and 25°C we have a resultant kinetic expression :

- (dXp/dt) = alXF a2(Xp/Xw)XF b2(Xp/Xw)XF2* (4)

DISCUSSION

The outstanding feature of the experimental results is the absence of a medium effect in the variation of the rate of decomposition of hydrogen peroxide with hydrogen peroxide concentration. The simplicity of kinetic form is helpful in eliminating from consideration as possible important factors in the reaction, properties of hydrogen peroxide + water mixtures which change in a complex way with composition. Thus the ionic product of the mixed solvent increases extremely rapidly, and eventually passes through a maximum, as the hydrogen peroxide concentration increases from zero. The concentration of H02- ions at constant pH will vary with the ionic product but this variation is not reflected in the rate of reaction. We can therefore say that the H02- ion is unlikely to be a kinetically important species in this reaction.

The rate of decomposition has been found to be proportional to Xp/Xw over a wide range of composition and this form of dependence strongly suggests that an important factor in determining the rate is a competition between hydrogen peroxide and water. It seems reasonable to suppose that this competition is for a place in the solvation sphere of the catalyst ion.

If we may represent a ferric ion in aqueous solution as an aquo-complex [Fe (H20)6]3+, then we may consider that, in hydrogen peroxide + water mixtures, an equilibrium of the form,

may be set up, in which a molecule of hydrogen peroxide competes with a water molecule for a place in the solvation shell of the ferric ion. Writing A for the aquo-complex [Fe(H20)6I3+, B for the mono-peroxy complex [Fe(H20)5H202]3+

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86 DECOMPOSITION OF HYDROGEN P E R O X I D E

and Y = Xp/Xw the mole ratio of H202 to H20, we may write for the equilibrium constant of this process :

Under our experimental conditions the oxidation of ferrous ion to the ferric state is virtually complete and we may put

whence

which becomes under the conditions 1 > K ~ Y

K1 = XB/xA Y.

XF = XA + XB,

X B = (KlXFY)/(1 + Klr), ( 5 )

XB = XFKlr ; (6)

(7) While this expression gives the correct form of dependence of rate of decomposi-

tion on peroxide concentration, it predicts that graphs of rate against Xp/Xw should pass through the origin, it contains no term in XF~, and it involves the assumption 1 > K1r even at very high values of X,. We may note that if K17” > 1 the expression predicts that the rate of reaction becomes independent of the peroxide concentration.

If we now suppose that a further substitution of hydrogen peroxide in the ferric ion solvation shell may occur and that both peroxy-complexes are catalytically active we have the scheme :

and supposing the rzite of decomposition to be proportional to X B we have - (dX,/dt) = k K1 XFY.

Ki [Fe(H20)6I3+ + M 2 0 2 + CFe(H2Q)s. H2O$+ + H2Q

[ F ~ ( € I ~ Q ) S W K W + + 11202 $ [Fe(H20)4. @I2Q2)2]3+ -t W2Q

[Fe@20)5 . H202]3+ 3 products,

[Fe(H20)4 . (H202)2]3+ 3 products. By a similar argument to that given above, writing C for the bi-proxy complex [Fe(H20)4 (H202)#+ and putting

we obtain XF = XA + XB + XC,

which under the conditions 1 < K1r > KlK21.2 reduces to

- (dX,/dz) = kl& + k2K2X~r. (9) This predicts that if - (dX,/dt) at constant XF is plotted against Y = Xp/Xw a straight line should result with a positive intercept on the rate axis. Furthermore the intercepts at different values of XF should be proportional to XF. These predictions are in accordance with the experimental results except that no term in X F ~ is obtained.

In order to account for the observed second-order term in XF, bimolecular reaction may be considered and an expression consistent with the observed behaviour is obtai2ed by supposing that the mono- and bi-peroxy complexes react together : B + C -$ products. This W O U ! ~ give a complete rate cxpressicn :

which under the conditions 1 < K1r > K1K2r2, reduces to

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P . JONES, R. KITCHING, M . L. TOBE AND w. F. K . WYNNE-JONES 87

Thus graphs of rate against r for fixed XF should be straight lines of slope (X~k2K2 f XF2k3&) and intercept ( k l X ~ ) . The intercepts of these graphs should be propor- tional to X F and (slope/XF) against XF graphs should be straight lines with positive intercepts on the ordinate. These features are in accordance with the experimental results at high peroxide concentrations.

At low concentrations of hydrogen peroxide the reaction of the mono-peroxy complex becomes dominant and the next region of the rate curve which may be treated simply is when the denominator assumes the form 1 + K1r > KlK2r-2. Then we may write,

(12) where the reciprocal of the rate becomes a linear function of l/r. kl may be obtained from the intercepts of these graphs and K1 from the ratio of intercept to slope. At sufficiently low peroxide concentrations this expression was approxi- mately obeyed and gave values kl - 20 moles peroxide min-1 moles catalyst-1 and K1 - 20.

At still lower values of X, the conditions (1 > K1r + K&r2) and X, --f 1 will be attained where

(1 3) We have not ourselves examined this range experimentally but it corresponds to the conditions used in previous studies of the reaction in dilute solution and the rate equation is in agreement with these earlier results. There is an evident need for more precise experimental work at intermediate peroxide concentrations, that is, the regions between the limits of the applicability of eqn. (13) and (11).

Although the ratio X J X , should strictly be replaced by the corresponding activity ratio we have confirmed, by using the activity data of Scatchard et aZ.10 that this is not a significant correction.

Eqn. (1 3) is in agreement with the results of previous work in dilute solutions of hydrogen peroxide, provided that the conditions Xp > X, applies. All the mechanisms which have been formulated for this reaction use the free-radical concept which was first applied to this particular reaction by Haber and Weiss.4 The catalyst ions are supposed to undergo alternative oxidation and reduction, forming in the process free radicals which are the active agents in bringing about decomposition of the hydrogen peroxide. It is generally agreed that an important step is the reaction,

which is sometimes written,

- l/(d&/dt) = (1/XFklKl) (l/r) f (l/XFkl),

- (dX,/dt) = ~ ~ X F X ~ .

Fe3+ + H202 -+ Fez+ + H02 + Hf,

Fe3+ + H02- + Fe2+ + HO2,

Fe2+ + H202 -+ Fe3+ + OH + OH-.

Since the catalyst is almost entirely in the ferric form, the rate of reaction (a) has the same form as the observed kinetics. One point of dispute is whether (a) is to be regarded as essentially rate-determining or whether, as is supposed by Barb et al.,11 it is the initiating reaction for a chain process which has the same kinetic form as reaction (a) under these concentration conditions. Barb et al. havc calculated the chain length on the basis of their proposed mechanism as the ratio of the overall rate constant to the rate constant (ka) for reaction (a). Unfortunately k, is not obtained directly but is expressed in terms of other quantities which are measurable according to their treatment. It may be noted that some of these quantities have been determined using ferric nitrate at an ionic strength of 0.435 and others using ferric perchlorate at an unspecified ionic strength, although Barb et al. give data which show that numerically different results are obtained using nitrates and perchlorates. Also their treatment gives, under different concentration

(a)

(a’)

(b’)

and it is also agreed that the ferric ion is regenerated in a reaction of the type,

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88 DECOMPOSITIOFJ OF HYDROGEN P E R O X I D E

conditions, a kinetic expression which is at variance with the observed results. Although these points do not necessarily invalidate their supposition of a chain process, the argument cannot be regarded as established. The present work provides no additional experimental evidence on this point since the reactions of the ferrous ion are extremely fast under our concentration conditions.

Barb et al. have considered the possibility that reaction (a) or (a’) may alter- natively be expressed :

Fe3+ + H02- 2 FeH022+ % Fez+ + H02, in which the unimolecular decomposition of the complex is slow. They have pointed out that this would change the order of reaction with respect to hydrogen peroxide to

that is, an order changing from first to zero as the fraction of the iron present as the complex increased. Barb et al. did not investigate conditions under which the validity of the different forms could be distinguished.

It is evident that the above treatment is similar to our scheme involving only a mono-peroxy complex and the resulting equations are identical provided that the change in reaction order occurs in the range where the conditions X, + 1 is applicable. However, the nature of the complex is by no means satisfactorily expressed by the equation of Barb et al. Their equation does not distinguish between a complex formed by displacement of water from the solvation sheath of the catalyst ion and the looser structure formed by the interaction of HO2- with a fully solvated ferric ion. A similar point concerning the dangers of describing reactions of this type, without consideration of the solvation of the ions concerned, has been made by Ubbelohde.12

We have given earlier reasons for supposing that the H02- ion is not an important species in this reaction. It seems preferable to regard the HO;! radical complex as being formed by the acid dissociation of a mono-peroxy iron complex followed by electron transfer :

[H2021/(1 + m 3 2 0 2 1 1 ,

[Fe3+(H20)5H202] + [Fe3+(H20)5H02-] + H+ .G

[Fez+ (H20)5H02] -+ products. A similar formulation is possible for the bi-peroxy complex. Re-oxidation of the ferrous ion could then occur via the analogous basic dissociation of a ferrous- peroxy complex :

[Fe2+ aq. H2,02] --f [Fez+ aq. HO+] + OH- + [Fe3+ aq. OH] -+ products.

Equilibrium in the first stage of this reaction sequence would imply exchange of oxygen between water and hydrogen peroxide in this system. Tracer experiments with 0 1 8 suggest that this is not the case although existing data are confined to dilute hydrogen peroxide solutions. 1 3 ~ 1 4

It then becomes possible to formulate processes leading to oxygen and water production which are entirely analogous to those suggested by Weiss and by Barb et al.11 In these it is supposed that the H02 radical becomes a free species and that oxygen is produced in encounters between H02 radicals and fsrric ions. Both the OH and H02 radicals may reasonably be supposed to propagate themselves through the solution by a series of fast hydrogen-atom transfer reactions with both com- ponents of the solvent. The endothermic reaction

is unlikely to be a participant in this scheme. A hydrogen-atom transfer reaction between an H02 radical and a hydrogen-peroxide molecule in the solvation sheath of a ferric ion would bring about the condition necessary for oxygen production.

HO + H202 + H202 + OH

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P. JONES, R . KITCHING, M. L. T O B E AND W. F . K . WYNNE-JONES 89

For the bi-peroxy complex an analogous intra-molecular process, which may be

[Fe3+ nq. 2H2021 + [Fe3f aq. H202H02-] + H+ written formally as :

[Fe3+ aq. H202H02-] + [Fez+ aq. HzOzH021

[Fez+ aq. H202H021 --f [Fe2+ aq. H02HO+] + OH- [Fe2+ aq. H02OH+] -+ [Fe3+ aq. H02HOI

[Fe3+ aq. H02HOI +- [Fe3+ aq. H20] + 0 2 .

becomes an additional possible reaction. It is possible to regard this reaction as analogous to a chain-terminating oxygen-producing step, between OH and H02, since if this path is important in the overall reaction, the type of chain visualized by Barb et al. becomes impossible.

As the concentration of catalyst is increased the possibility of reaction via the encounter of peroxy-complexes arises. We have shown that the kinetic term in X F ~ could arise in this way by the interaction of mono-and bi-peroxy complexes. This may be imagined as occurring by the reaction sequence :

[Fe2+ aq. HO2] + [Fez+ aq. H202H021 =+ [Fez+ aq. H2021 + [Fe2+ aq. H202. 0 2 1

[Fe2+ aq. H202 , 0 2 1 -+ [Fez+ aq. H2021 + 0 2 .

A reaction according to this scheme would show no primary salt effect since the pre-equilibrium is a homolytic atom transfer and neither step involves a change of charge distribution. This is in accordance with the experimental results. A similar re- action between two mono-peroxy complexes is also possible but our kinetic evidence suggests that this is not important at 25"C, although it may become important at elevated temperatures. It is possible to regard these bimolecular reactions as resulting in the formation of an unstable binuclear molecular oxygen complex. There is considerable evidence in the literature which suggests this may well be the case. Binuclear cobalt complexes of this type have been shown to exist by Werner 15 and a complex of this type is considered by Baxendale and Wells 16 to be the active entity in the reduction of cobaltic salts by hydrogen peroxide. The results of Lamb and Elder 17 and of George 18 on the autoxidation of ferrous salts in aqueous solution show the rate to be proportional to the square of the ferrous ion concentra- tion and directly proportional to the oxygen pressure. This suggests the possibility of the formation of [Fe2+ 0 2 . Fe2+] as an intermediate, but the kinetic evidence will not necessarily distinguish this from a scheme involving the formation of, say, Fe2+ . 0 2 followed by a reaction of the type :

Fe022f + FeH202+ + Fe2fH02 + Fe2+HO. Since, in the reaction scheme we have proposed, the concentrations of the active

species depend upon the acid-dissociation of the peroxy-complexes, the rate of reaction at constant peroxide concentration would be expected to decrease as the pH of the solution decreases. This is generally confirmed by our results but further work at constant ionic strength is required.

Since our kinetic evidence implies that the rate constants are independent of concentration it must be concluded that the peroxy-complexes are sufficiently strong acids for their dissociations to be insensitive to the medium change. The acid dissociation constant of [Fe(H20)6]3+ is of the order 10-2. The acidity of hydrogen peroxide is about 106 times that of water and if we assume a comparable acidity ratio in the ferric aquo- and peroxy complexes then the peroxy complexes would be expected to be reasonably strong acids. This formulation avoids the difficulty encountered in supposing that the pH dependence of the reaction rate is explained by writing a reaction such as

Fe3+ + 02H- + Fe2+ + HO2,

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90 DECOMPOSITION OF HYDROGEN PEROXIDE

the rate of which would change in a complex way with hydrogen peroxide concentra- tion since it responds to changes in the ionization constant of the mixed solvent.

It is necessary to consider whether any direct experimental evidence exists for the peroxy-complexes formulated. Solutions of iron salts in concentrated hydrogen peroxide have a deep brown colour and Evans, George and Uri 19 ascribed this colour to the electron transfer process,

Fe3+HO2- -+ Fe2+H02,

but were only able to interpret their experimental results up to hydrogen peroxide concentrations of about 10 M. At higher concentrations the optical density increased rapidly and this was ascribed to a " superacidity " effect, which it was supposed might result in the formation of a ferric perhydrate. A recent analysis 20 has shown that the data given by Evans et al. (their fig. (2), referring to cHC104 = 0.5 M and hydrogen peroxide concentrations up to 33 M) are consistent with the equation given above deriving the concentration of the ferric mono-peroxy complex. Further experimental work has verified this conclusion. Preliminary experiments at pH = 2 have given results which are consistent with the formation of both mono- and bi-peroxy complexes. It is hoped to extend the spectrophotometric studies to obtain more quantitative information on the peroxy-complexes of iron.

1 Schumb, Satterfleld and Wentworth, Hydrogen Peroxide (Reinkold, New York, 1955). 2 Baxendale, Advances in Catalysis (Academic Press, New York, 1952), 4, 31. 3 Weiss, Advances in Catalysis (Academic Press, New York, 1952), p. 343. 4 Haber and Weiss, Proc. Roy. SOC. A, 1934,332,1473 ; Weiss, Faraday SOC. Discussions,

5 Kilde and Wynne-Jones, Trans. Faraday Soc., 1949, 47, 616. 6 Banfield and Hilden, Laporte Chemicals, private communication. 7ref. (l), p. 394. 8 Mitchell and Wynne-Jones, Trans. Faraday SOC., 1955, 51, 1690. 9 Bell and Baughan, Proc. Roy. SOC. A , 1937, 158, 464 ; Cullis and Peard, Trans.

10 Scatchard, Kavanagh and Ticknor, J. Amer. Chem. Soc., 1952,74, 3715. 11 Barb, Baxendale, George and Hargrave, Trans. Faraday Soc., 1951,47,462 and 591. 12 Ubbelohde, Favaday SOC. Discussions, 1947, 2, 215. 13 Bunton and Llewellyn, Research, 1952, 5, 142. 14 Cahill and Taube, J. Amer. Chern. SOC., 1952,74,2312. 15 Werner and Mylius, 2. anorg. Chem., 1898, 16,245 ; Werner, N. Ansch., 1923, 274. 16 Baxendale and Wells, Trans. Faraday SOC., 1957, 53, 800. 17 Lamb and Elders, J. Amet. Chem. SOC., 1931, 53, 137. 18 George, J. Chem. SOC., 1954,4349. 19 Evans, George and Uri, Trans. Faraday SOC., 1949,45, 230. 20 Jones, in preparation.

1947, 2,212.

Faraday SOC., 1949, 47, 616.

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