THE RADIATION CHEMISTRY OF FERROUS-FERRIC SYSTEMS PART 1. REACTIONS IN AIR-EQUILIBRATED SOLUTIONS BY C . B. AMPHLETT Chemistry Division, A. E. R. E., Harwell, Berks. Received 14th February, 1952 The oxidation of ferrous ions by X- and y-radiation has been studied in dilute solu- tions in &SO4 at different acidities. The kinetic scheme proposed to explain the resuits in 0.8N H2S04 has been found inadequate to explain the variation in initial oxidation yield with pH ; it is suggested that the primary act is more complex than is usually assumed. From a study of the steady-state condition at different acidities, values have been ob- tained for the “ equivalent redox potential ” of irradiated water with respect to the ferrous- ferric system. The effect of addition of other ions and complexing agents is briefly described.The oxidation of dilute aqueous solutions of ferrous ion in acid solution was first studied by Fricke and his co-workers,l who suggested it as a means of measur- ing the integral energy absorption in irradiated solutions. In 0-8 N H2S04 oxidation is virtually complete and ferric ion hydrolysis negligible, while the high yield makes it possible to measure moderately low doses with the sensitive analytical met hods available .2 The following mechanism has been proposed3 to describe the oxidation in the initial stages in sufficiently acid media (> 10-1 N). H20 y2- H + OH Fe2f + OH --f Fe3+ + OH- Fe2f + €402 --f Fe3+ + H02- H02- + H + Z-f H202 Fez+ + H202 -+ Fe3f + OH- +- OH HA- + OH- ~2 H2O. The maximum theoretical ratio between the oxidation yields in aerated and evacuated solutions is Ii = 4/1, whereas in practice a figure of about 2.7 is ob- tained.(The value of 4 reported by Krenz and Dewhurst 3 has since been traced to contamination of the evacuated samples, possibly with mercury vapour, leading to a lowered evacuated yield (Dewhurst, private communication).) Hart has suggested 4 that a second “ primary ” step be added, viz., This represents the increased probability of pairwise combination of like radicals in regions of dense ionization (e.g. along a-particle tracks, or near the end of /3-particle tracks), where the local radical concentrations are high. It had pre- viously been postulated by Allen 5 to explain production of H202 in “ hot spots ” in irradiated water.It should not be confused with similar combination re- actions in the bulk of the solution ; the results of ferrous 4- H2Oz kinetic studies 144C . B. AMPHLETT 145 at niuch higher rates of radical production than those in this work show that the OH -t OH reaction can be neglected in comparison with step (l), while the linearity of yield up to complete oxygen depletion in irradiation experiments shows that the H + H reaction is negligible compared with step (2). By adding (0’) to the above sequence, and assigning to the two steps probabilities x and (1 - x), the aerated yield is decreased without affecting the evacuated yield, thus reducing R to less than four. By considering the ratios experimentally obtained, Hart finds that the importance of (0’) increases with increasing ion-density (e.g.from Co y-radiation to T3 P-radiation), in accordance with expectation. It is found, however, that the ratio of x/(l - x) appears to depend upon the solute, being appreciably different for dilute formic acid solutions.6 This would suggest that Hart’s treatment is an approximation, which it is convenient to employ in our present state of knowledge concerning the primary act. In 0.8 N H2S04 the oxidation yield is found to be independent of initial ferrous ion concentration over a wide range, viz., 10-4 to 10-1 M.7 The decrease in yield below 10-4 M is attributed to the relative slowness of step (4) under these con- ditions. If sufficient time is allowed between irradiation and analysis, normal yields are obtained,g and it has been shown 9 that the rate constant for the increase in yield on standing is in agreement with that for step (4).In 0.8 N H2S04 the yield is constant almost to complete oxidation, and oxidation is virtually complete, but as the pH is increased the initial yield eventu- ally decreases and evidence of a back reaction appears; oxidation is no longer complete, and the oxidation curve bends over to a steady state with respect to ferrous and ferric ion. In addition, reduction of aqueous ferric solutions, which is barely detectable in strong acid, is observed to an increasing extent. These effects may be attributed to the reduction of ferric ions, which are known to com- pete for HOa radicals in the pH-dependent step Baxendale and his co-workers have shown10 in a study of the Fez+ -I- Hz02 reaction, which has many features in common with the present system, that the ratio kS/k3 (k4/k3 in their terminology) is pH dependent, and that the reacting partners in step ( 5 ) are probably Fe3f and 02-.The possibility of ferric ion being reduced by the step Fe3+ +- HO2- + Fez+ + H02, i.e. the reverse of (3), is ruled out by its extremely low rate compared with the ferrous oxidation steps, so that it would only become appreciable at very high ratios of ferric to ferrous ion. Little work has hitherto been reported on the irradiation reaction at low acidities,ll possibly because of the difficulties introduced by hydrolysis of ferric ion; it has been shown 12 that as the pH is increased a progressive lowering of the steady-state [Fe3+]/[Fe2+] ratio is obtained.It is the purpose of this paper to discuss the effect of pH in air-equilibrated solutions, and also to detail some of the effects observed in the presence of complexing agents and added ions. Fe3’ . j HOz -tFe2+ -i- Hi 4- 0 2 . ; . ( 5 ) EXPERIMENTAL MATERIALS.-h the original experiments 12 perchlorates and perchloric acid were used to avoid complexing other than that arising from hydrolysis. It was found, how- ever, that below pH 2 the yields were appreciably higher than in sulphate solutions ; in addition chlorate (but not chloride) was formed. Weiss has since found13 a similar phenomenon in air-free solutions, together with the appearance of chloride. It appears that perchlorate ion must participate in the reaction, and sulphate solutions have since been used, despite their great complexing power.Nitrate ion was excluded, because of possible complications arising out of its reduction to nitrite ; 14 chloride ion is known to compete for OH radicals.15 A.R. ferrous ammonium sulphate was twice recrystallized from 0.8 N HzSO4 and dried in air ; all other materials were A.R., and were not purified further. Distilled water was redistilled before use, once from 0.01 M KMn04 + 0-01 M KOH, once from 0.01 M KHS04 and finally alone. All distillations were from silica apparatus. and the final146 FERROUS-FERRIC SYSTEMS water was stored in stoppered silica flasks. This water was used for all irradiated sohi- tions, and for the final rinsing of apparatus. (i) 400 mc Ra y-source in Pt capsule. IRRADIATIONS.-Three sources have been used, viz.: Samples were irradiated in Pyrex vessels mounted coaxially with respect to the source. Dose rate 33 r/min (1.9 x 1015 eV/cm3 min). Samples irradiated either in polystyrene cell containing 10 ml solution (208 r/min ; 1.22 x 1016 eV/cm3 min) or in 2 cm3 glass cells mounted in a polystyrene block (338 r/niin ; 1.99 x 1016 eV/cm3 min). 15 cm3 samples were irradiated in thin-windowed Pyrex cells mounted close to the shutter. Dose rate 1420 r/min (8.4 x 1016 eV/cm3 min). All dose rates were measured by dosimetry with 0-8 N HzSO4 solutions of ferrous ion (2 x 10-4 M to 2 x 10-3 M), using a value of G = 20.6 ions oxidized per 100 eV.2 ANALYsIS.-After irradiation the solutions were analyzed for Fez+ by complexing with a-phenanthroline at pH 4-5 in an acetate buffer,16 measuring the optical density at 510 mp with a Hilger Uvispeck spectrophotometer.The calibration curve for ferrous ion gave a value of eM = 10,660 for Feph$+ under these conditions (ph E a-phehanthrol- ine). This is low compared with the accepted value of 11,000-11,100, but was consistently obtained with this instrument and the particular cells used. The optical density of a ferric solution (free from Fez+) containing o-phenanthroline and acetate corresponded to eM -20, so that the contribution of ferric ion species towards the optical density is negligible. pH values were measured with a glass electrode and Beckman pH meter. (ii) 3.25 curie Co y-source in stainless steel container. (iii) 240 kVp X-rays, unfiltered, at 8 mA, from a Victor Maximar set.1. THE INITIAL OXIDATION YIELD AS A FUNCTION OF ACIDJTY RESULTS Table 1 shows the results obtained in H2SO4 solutions from 0-80 N to 1.6 x '10-4 N in H+ ion; table 2 gives for comparison some earlier results in perchlorate solutions. The absolute accuracy of the sulphate figures is probably :t 5 % ; the perchlorate results are less accurate, as they were obtained before conditions had been thoroughly standardized. TABLE ~.--JNITIAL OXIDATION YIELD OF FERROUS ION IN H2SO4, IRRADIATED WITH [Fe2+]o = 1-5-2.0 x 10-4 M. Dose rate 33-1420 rjmin =1-9 x 101s - 8.4 x 1016 eV/cmJ min. x- AND y-RADIATION [H2S04], moles/l. PH [H+l moIes/I. Go = no. of ions oxidized per 100 eV 0.4 (by titration) 0.39 (calc.) 0.41 20-6 0.05 ,, 1.21 6.0 X 10-2 20.6 -2.0 x 10-2 1.60 2.51 x 10-2 20.6 -1.0 x 10-3 2.65 2-24 x 10-3 16-2 -3.6 x 10-4 3-10 7.9 x 10-4 13.9 -1.3 x 10-4 3.518 2.6 x 10-4 10.3 -9.0 x 10-5 3.80 1.6 x 10-4 8.4 1.0 TABLE 2.-INITIAL OXIDATION YIELDS IN PERCHLORATE SOLUTIONS 150-180 kVp X-rays; dose rate 270-370 r/min =- 1.59-2.18 :: 1016 eV/cm3 min.PH [H+], moles/l. Gn 0.1 (calc.) 0.80 41-2 1 -05 9-8 x 10-2 29.0 2-06 8.7 x 10-3 18.1 2-20 6.3 18.1 3-00 1.0 14-5 3-10 8.0 x 10-4 12.5 3-64 2.3 12-5C . B . AMPHLETT 147 DISCUSSION The variation in Go with pH is shown in fig. 1, which also shows results ob- tained by other workers. It is not proposed to consider the perchlorate results beyond remarking that above pH 2 there appears to be no difference between perchlorate and sulphate solutions ; it may reasonably be conjectured that in this region we are dealing principally with uncomplexed ferrous ion and the ion FeOH2+ in both cases.The kinetics outlined in the introduction give a qualitative explanation of the oxidation process, and in strongly acid solutions conform to the observed stoichio- metry. If we assume that the distribution of atoms and radicals arising from y-irradiation is homogeneous, we may apply steady state kinetics to obtain an FIG. 1.-Variation in initial oxidation yield with pH. 0 SO$- solutions (this work) 6 ,, ,, (Krenz and Dewhurst, unpublished results) A 3 , ,, ; 10-3 N H2S04, 0.3 M Na2S04 El c104- ,, expression for the yield. From the calculated steady state concentrations of H, OH, H02 and H202, the overall expression for the yield of ferrous ions oxidized per 100 eV may be calculated to be where D is the dose rate in units of 100 eV per unit time, ko is the net radical pair yield per 100 eV (allowing for the back-reaction in the primary step) and x is the fraction of radicals available for reaction with substrate in the bulk of the solution.The first term in parentheses represents the stoichiometry 1 OH = 1 Fe2+, while the second represents competition between ferrous and148 FERROUS-FERRIC SYSTEMS ferric ion for the H02 radicals, each of which is capable of oxidizing three ferrous ions. In the initial stages of the reaction, when ks[Fe34] < k3[Fe2+], this reduces to the form (2) In evacuated solutions Go' = ko, and hence x may be evaluated. Using a value of R = Go/Go' = 2-7-2-86, we obtain x = 0.57-0.62, say 0.6.This implies a net radical pair production of slightly more than seven per 100 eV, corres- ponding to an energy requirement of about 13 eV per radical pair in liquid water. Although the initial portion of the oxidation curve remains linear as the pH is increased, the initial yield decreases above pH 2. The form of the curve re- lating Go to pH is different from that given by Fricke and Hart,ll the reason for this discrepancy not being obvious. On the basis of the above mechanism, the decrease in yield must be due to a decrease in concentration of one or more of the reactants concerned, viz. Fez+, OH, HO;! and H202. Since the decrease is apparent from the start of the irradiation, and since the curvature due to back reactions involving ferric ion is not apparent until much later, we cannot invoke ferric ion effects to explain the decrease in initial yield.The hydrolysis of ferrous ion is negligible at these acidities ( K h = 1.2 x 10-6 17), and it would appear that sulphate complexing has no effect; the effect of adding excess neutral sulphate to a solution at pH 3 is to depress the yield to a value corresponding to the increase in pH. It appears unlikely that the variation in initial yield can be attributed to differences in reactivity between hydrated ferrous ions and possible complexes. The results of other workers on the ferrous + H202 reaction 10 show that kq, deter- mined in HC104 and H2SO4 solutions under conditions where steps (1) and (4) are predominant, does not vary over the range N/2 acid to pH 2.65 ; this confirms the conclusions drawn from the present work.The oxidizing power of the irradiated solvent must therefore vary, and there are several alternatives. The dissociation constant of H202 is so low (Kd = 1-8 x 10-12 18) that we can eliminate the possibility of step (4) being pH dependent at moderate acidities. Similarly, the dissociation of OH into 0- and H+, followed by 0- + H --f OH- would appear unlikely; 19 both these steps are, of course, eliminated also by the results of the Fe2f + H202 studies. Weiss has suggested 20 that in air-free solutions the following pH dependent steps may operate : Go =: ko(1 + 3 4 . H + H+ z-f H2+ H2f + Fez+ -+ H2 + Fe3+. In the presence of 0 2 the equilibrium would be expected to be well on the side of H atoms, and such a mechanism in aerated solutions would actually lead to a decrease in yield with increasing acidity. The participation of OH+ ions,21 and their removal by OH+ + OH- -+ H202 will have no effect in the ferrous system, since both OH+ and H202 are equivalent to two Fe2+ ions.This leaves only H02 as a source of pH dependence. It has been pointed out 10 that whereas H02 behaves as an oxidizing agent towards ferrous ions, its anion 0 2 - will reduce ferric ions, so that with decreasing acidity the net oxidizing power should be reduced. K d for the dissociation HO2 z-f H+ + 0 2 - has been estimated as 10-3, with a possible error of about one power of ten in the value.22 If we modify eqn. (2) to allow for the fraction of H02 radicals dissociated, and therefore whenceC .B . AMPHLETT 149 Gmax, the maximum yield when Kh < [H+J will be given by (2), so that Substituting for x, Kd and Gmax, we obtain a curve relating Go and pH of the form shown in fig. 1. It will be seen that, although the decrease in yield is found to be in the correct acidity range, the forms of the theoretical and experimental curves are in fact appreciably different. Thus, while the dissociation of HO2 is probably contributory, it cannot be the only cause. The experimental values between pH 2 and 4 were found to fit the empirical relationship which is also plotted in fig. 1 ; below pH 2 the values deviate appreciably from this curve. The complexity of this expression suggests a complicated overall pH dependence, with possibly several factors operating.It should be remembered that ko and x may not be independent of acidity, as has been assumed in this treatment, and that the primary act of radical production may itself be much more complex than has hitherto been assumed. Lefort has suggested23 that an addi- tional source of primary radicals may arise from the pH dependent sequence H30f + e --f H2 4- OH (6) H2 -k OH -+ HzO + H. (7) If these steps be added to the primary step (0), it is no longer necessary to assume the production of equal numbers of H atoms and OH radicals, and the production of both species is pH dependent. However, such a step will introduce a pH dependence in opposite direction from that found in ferrous solutions, since increasing acidity will favour capture of the slow electrons in step (6) rather than by the process H20 + e -+ H20- -+ H + OH-, thus leading (in aerated solu- tions) to a decrease in oxidizing power.It seems therefore that a more complete understanding of the primary act in irradiated water is necessary for interpretation of the overall pH dependence ; further work is proceeding with evacuated solutions in an attempt to obtain more information. 2. THE STEADY STATE AND THE EQUIVALENT REDOX POTENTIAL OF IRRADIATED WATER On the above mechanism the value of G at any point on the oxidation curve is given by the expression (1). From the study of the Fe2f 4- H202 reaction, values of k5/k3 may be calculated ; 24 they range from 1-25 at pH 4 to about 0.02 in 0.8 N acid at 25" C. It is found in irradiation experiments that the point at which the yield begins to deviate from its initial value Go approaches more closely to the origin as the pH is increased, due to steadily increasing values of k5/k3.In 0.8 N H2SO4 it is so near to 100 % oxidation as to be frequently missed ; if, however, a large excess of ferric ion is initially present, thus increasing [Fe3+]/[Fe2+], the deviation from linearity is easily observed, even though the initial yield is unchanged.8 An attempt to apply expression (I) quantitatively fails for two reasons : (i) we should expect a continuously decreasing yield as oxidation proceeds, at acidities where k5/k3 is appreciable. As fig. 2 shows, the initial yield is constant for an appreciable part of the oxidation and this is still true at pH 3.8 ; (ii) the expression (1) will never yield a steady state value of G = 0, except at almost complete oxidation, for positive values of kS/k3 and of x.This suggests the introduction of some other reducing step such as (8) Fe3+ -I- H -+ Fez+ + H 1 .150 FERROUS-FERRIC SYSTEMS This will itself presumably be pH dependent, and over practically the whole range of acidities used in these experiments the contribution of FeOH2+ cannot be neglected. The substitution of step (8) into the original kinetic scheme does not, however, solve the difficulty, since we obtain negative values of ka[Fe3+]/k2[021 on substituting for kS/k3 and [Fe3+]/[Fe2'] when G = 0. It is clear, therefore, that a fundamental inadequacy of the kinetics has been exposed. Although a complete elucidation of the kinetics is not possible at present, the results obtained for the steady state ratio of ferric to ferrous ion may be used to calculate the " equivalent redox potential " (e.r.p.) of irradiated water.The concept of e.r.p. has been fully discussed by Collinson and Dainton,25 who con- clude from a survey of results on the irradiation of aqueous solutions of oxidizing and reducing agents that for X- and y-irradiation e.r.p. (H20) is somewhat morc positive than - I-OV (U.S. convention). A few examples of steady state con- &ions attained in irradiations are known, e.g. the systems I NO. --NO3 ,2b FIG 2.-Oxidation of ferroug solutions to steady state conditions at pH 391. 240 kVp X-rays, 1400 r/min Co y-rays (1.2 mV) 208 r/min Ra 'lV mean)7 33 r/min x 1-46 x 10-4 M Fe2+ 0 1.46 X 10-4 M Fez i- 0 1-48 ,, 6 1-54 ,, 7, t 1.52 ,, 6 1.66 ,, 7 , If 1-72 ,, Q 1.72 y y Y 7 0 1-95 ,, 7 7 -0- 1.82 ,, , I 0 2.04 ), 7 , The figures to the right of the curves refer to the percentage oxidation at the steady state.The horizontal scale has been considerably compressed above 10 x 1017 eV/cm3. T2-I-,26 arsenite--arsenate,29 and Fe*+-Fe3+ ; in these cases the steady state may be approached from either side, and no effect is observed if one irradiates a solution of composition corresponding to the steady state. From the steady state value of [N03-]/[N02-] = 1.5 quoted by Lefort26 and from the known EO value for this system (- 0.94 V 27 in acid solution, corrected to - 0-84 V at pH 6), we may calculate a value of - 0.49 V for the e.r.p. of X-irradiated water assuming the overall reaction to be H20 + NO2- ,f NO3- -I- 2HI + 2e.2-04 x 10-4 M Fe2tC. B . A‘MHPLETT 151 RESULTS We have measured steady state values in the Fez+ + Fe3 1- system in H2SO4 at different acidities, both for ferrous oxidation and for ferric reduction. Fig. 2 shows the results at pH 3.1 for X- and y-radiation over a forty-fold variation in dose rate. No variation either in initial yield or in steady state conditions was observed over this range. In order to calculate values of e.r.p. (H20) we have measured Eo’ values for the ferrous + ferric system as a function of pH from 0 to 3-5, using a Cambridge millivoltmeter and a saturated calomel electrode. The variation is complex, and is being studied further. At high acidity our results are in agreement with other workers,28 but no other results are known to us for lower acidities. Using these values of Eo’ we have calculated e.r.p.(HzO) from the steady state ratios by means of the Nernst formula e.r.p. (H20) = Eo’ - 0.059 loglo [Fe3 ‘]/[Fe2+]. We use the symbol Eo’ to denote the observed redox potential relative to the hydrogen electrode, under the conditions of measurement, when [FelI~l = [Fe“] ; this should not be confused with the standard thermodynamic potential EO- Table 3 summarizes the results so obtained ; the absolute values of e.r.p. (H20) are probably correct to 1 5 mV. TABLE 3 .----EQUIVALENT REDOX Po rEmiAL OF IRRADIATED WATER IN 7 HE system PI1 Fez+ 2.12- oxidation 2.1 5 247 3.10- 3.1 5 Fe3+ 2.12 reduction 3.22 Fe3 + reduction 2.08 0 3 2 - saturated) radiation Y Y Y X x Y X FERROUS-FERRIC SYSTEM [ Fe3+]/ [Fe2+] at steady state Eo’ 6.20 --0*704 V 6.30 5.30 3.28 -0.708 3-10 3-1 1 3.67 -0.703 3.25 4-14 3-06 4.22 3-76 7.1-9.3 -0.704 3.28 -- 0.702 4.94 -0.703 e.r.p.(H20) mean e.r.p. (H20) - 0.750 V 0.75 1 - 0.749 &0*002 V 0.747 - 0.738 0.737 0.737 -0*737-J:0.001 - 0.736 0.733 0.739 - 0.736 1 0.004 0.732 0.740 0-737 -0.754 to -0*757+0.003 - 0.760 -0.732 - 0.732 - 0.744 -- 0.744 J l L % C J L S $ u T r J As the pH is increased, the value of e.r.p. (H20) becomes more positive, in conformity with the general trend of redox potentials with respect to increasing alkalinity; the change in oxidizing power thus parallels the change in initial oxidation yield. The presence of H2, which is assumed to remove OH radicals by the reaction H2 + OH -+ H20 + H, produces only a slight change in potential a t pH 2. We can calculate values of e.r.p.(H2O) for 0.8 N and 0.1 N acid, putting [Fe3+]/[Fe2+] ’v 100 ; we then obtain values of N - 0-806 V in 0.8 N acid, and N - - 0.799 in 0.1 N acid, showing an appreciable decrease in e.r.p. to more negative values below pH 2. These values differ appreciably from that calculated from the results of Lefort’s work on the system N02--N03-; this may in part be due to differences in pH,152 FERROUS-FERRIC SYSTEMS but may also be due to the different substrates present in the two cases. The measured e.r.p. of water relative to any dissolved material will depend upon the competition for the oxidizing and reducing species formed between the substrate (in both oxidized and reduced forms), the solvent and the radicals themselves; in the presence of dissolved oxygen the situation will be still more complicated.It is therefore conceivable that the precise value of e.r.p. will depend upon the particular substrate chosen, as well as upon other external conditions, such as acidity, temperature, presence of oxygen, hydrogen, etc. Thus, although in principle we can conceive the idea of irradiated water possessing a redox potential, it may not be possible to measure it precisely by experiments of this kind. All the possible reactions involved in the oxidation and reduction of ferrous and ferric ions in aqueous solution are appreciably exothermic, with one ex- ception ; their AH values, calculated from cycles involving the entities concerned, are given below: --f Fez+ + H-+ ; - 63.6 ,, (iii) - 16.0 ,, (vi> The reduction of hydrated ferric ions by hydrogen atoms is energetically favourable, but in view of the pH dependence it seems more likely that reduction involves the hydrolyzed species FeOI-l2+.Work on the photo-decomposition of ceric solutions 30 and on the radiation induced reduction of acid ceric solutions 31 suggests that the principal reduction step is CeOH3t t OH + Ce3+ + H202. This is energetically more favourable (AH = - 20.1 kcal30) than (v) above ; in addition, the peroxide produced reduces more ceric ion, whereas in the present system it would oxidize ferrous ion and so produce no net reduction. The precise function of H atoms in producing reduction in aqueous solutions seems at present to be doubtful; although it appears likely that they participate in the ferric system, recombination to molecular hydrogen seems to predominate in the reduc- tion of ceric solutions.32 It is hoped to obtain further information by studying the H202 yield and gas production in evacuated systems.The apparent anomaly in the effects of acidity upon the FeZ’-Fe31- and Ce3+-Ce4+ systems disappears when viewed in the light of the potentials in- voked. In acid solutions, oxidation of ferrous solutions is predominant, while oxidation of cerous solutions is only observed in alkaline solutions.31 The above measurements show that e.r.p. (H2O) becomes slightly more positive with in- creasing pH, while the potential of the CelI1/CetV couple changes from - 1.44 in acid solution to + 0.77 V over the range of conditions employed by Lefort.31 Consequently the small change in e.r.p.is overtaken by the change in Eo’. There are no results for ceric solutions over the range of acidity used in the present work, owing to difficulties introduced by the hydrolysis of ceric ion. Under the con- ditions employed for cerous oxidation, viz. in solution in 3 M K2CO3, the cerous ion will be largely present as anionic complexes and under similar conditions it is probable that ferrous solutions would again be oxidized ( E l = + 0.56 V). It appears from the present work that although oxidation of hydrated ferrous ions is easily affected, reduction is greatly favoured by hydrolysis of ferric ion; this may be due to a greater ease of electron-transfer between the reducing species and the central ion when the hydration shell is broken, either by hydrolysis or by suitable complexing.Silverman and Dodson 33 have reported that decrease in acidity favours isotopic exchange between ferrous and ferric ions in perchlorate solutions, and attribute this to a greater facility for electron-transfer in the inter-C. B . AMPHLETT 153 mediate state Fe . . . OH- . . . Fe compared with that in Fe . . . H20 . . . Fe. The experiments on I‘ optical interaction absorption ” in mixed ferrou?+ ferric systems 34 are also taken to indicate increased probability of electron-transfer when the hydration shells are broken. It is known that ceric solutions are much more strongly hydrolyzed than the corresponding ferric solutions (92 % 30 in N HC104, compared with 0.5 % 3 9 , while a comparison of the respective EO values in HC104 and H2SO4 solutions27928 suggests that ceric ions are much more strongly complexed. Further information should be sought on the effects of pH on the irradiation of redox systems, since it may lead to knowledge con- cerning the reactivities of ions and complexes in solution.3. THE EFFECT OF ADDED MATERIALS ON THE YIELD AND STEADY STATE VALUES FLUORIDE ioN.-Fluoride (0.01 M) added to solutions with a final pH of 2.8 and 4.2, produced practically complete oxidation, with no evidence of a ferric back-reaction ; complete oxidation was still observed in 0.8 N HzSO4 solutions. This agrees with the change in redox potential when fluoride is added, e.g.addition of F- to a 0.1 N HC1 solution displaces Eo’ from - 0.771 V to -- 0.553 V at N/70 F- ion concentration.36 Addition of fluoride also affects the initial yield in a com- plex manner. Thiocyanate, which also complexes strongly with ferric ion, produces similar results to fluoride at high pH. CHLORIDE ioN.-Chloride ion had little effect on the steady state ratios. Contrary to the results reported by Dewhurst,37 a marked effect was found upon the initial yield in low chloride concentrations, the diminution in yield increasing with increasing C1- ion concentration ; values obtained from several runs were Go = 17.2 (lO--3 M Cl-), 15.2 (10-2 M Cl-), 9.1 (M Cl-), all in 0.8 N H2SO4, and 8.3 (10-2 M and 1 M) at pH 3.1. The reason for the discrepancy is not obvious, but it is proposed to examine it further in air-free solutions. A~RYLoNITR~LE.-A ddition of many reactive organic compounds has been shown to increase the ferrous ion consumption both in the irradiation reaction 37 and the Fenton reaction.38 On the other hand, addition of acrylonitrile (1 M) to a 0.8 N H2SO4 produced a decrease in yield (Go -x 13.8) without affecting the overall oxidation ; the latter point is to be expected, since only a 1 mV change in Eo’ is produced.39 Acetonitrile (1 M) produced a similar effect (Go = 12.2).In these cases the effect may be due to competitive attack by the radicals upon the unsaturated structures present, presumably producing organic radicals which are not capable of peroxidizing and so causing an induced oxidation of ferrous ion.PHENANTHROLINE AND DIPYRIDYL-The most striking effect was observed in the case of the o-phenanthroline and dipyridyl complexes, where complete re- duction of the ferric complex was observed from pH 1 to 4, with no oxidation of the ferrous complex.12 These form a reversible redox system (EO = - 1-10 V). At low pH the initial yield is quite high (Go _N 12 at pH 1.0), suggesting great ease of electron transfer for such bulky ions; this is supported by the results of radioactive exchange experiments.40 At pH 2 to 3, GO = 4 to 5, similar to the value found at pH 6 in the absence of air ; 31 we also find that the yield in this region is unaffected by 0 2 , N2 or H2. It was found during the original investiga- tions that the ferric complex was also reduced by H202; 41 the kinetics of this reduction, as well as those of the irradiation reaction, are being further studied.The author wishes to thank Sir John Cockcroft, F.R.S., for permission to publish this paper; his grateful thanks are due to Mrs. M. 0. Small for con- siderable assistance with the experimental work, to Dr. W. Wild for discussion of the results, and to the M.R.C. unit at Harwell for the loan of facilities for X-irradiation.154 PEKKOUS-FERRIC SYSTEMS NOTE ADDED IN PRooF.-Since writing my paper, further evidence has appeared concernjng the relative contributions of steps (0) and (0‘) in the mechanism dis- cussed. Measurement of the moiecular Hz yield in aerated ferrous solutions in 0.8N H2SO4 suggests 42 that x := 0.78, and not 0.62 as was estimated from Hart’s earlier figures.This assumes that thc yield of H202 in the “hot spots ” equals the molecular H2 yield, which is supported by the measurements of Sutton 43 on the post-irradiation reaction in very dilute (< 10-4 M) deaerated ferrous solu- tions. On substituting the above value of x the expression (4) is altered slightly, the factor 3x/(1 + 3x) changing from 0.645 to 0.70 ; this will not, however, greatly affect the position of this curbe in fig. 1. The discrepancy between this value of x and the earlier one derived by Hart is readily understood if we consider the mechan- ism for the oxidation of deaeratcd solutions of ferrous ion.44 The participation of H2-‘ in the latter case, but not in aerated solutions, implies that calculations of x based solely on air/vacuum yield ratios may be considerably in error, par- ticularly since the ratio is dependent upon ferrous ion concentration.1 Fricke, Physic. Review, 1928, 31, 1117. 2 Miller, J. Chem. Physics, 1950, 18, 79 ; also this Discussion. 3 Krenz and Dewhurst, J . Chem. Physics, 1949, 17, 1337. 4Hart, J. Amer. Chem. SOC., 1951, 73, 1891. 5 Allen, J . Physic. Chem., 1947, 52, 489. 6 Hart, ANL-4434 (U.S. Atomic Energy Commission, declassified document). 7 Miller, ref. (2). Todd and Whitcher AECU-458 (US. Atomic Energy Commission, declassified document). 8 Krenz, unpublished work. 9 Dainton and Sutton, unpublished work. Fricke and Morse, Phil. Mag., 1929, 7, 129. 10 Barb, Baxendale, George and Hargrave, Trans. Furatlay Soc., 195 1, 47, 462. 11 Fricke and Hart, J. Chem. Physics, 1935, 3, 60. 12 Amphlett, Nature, 1950, 165, 977. 13 Weiss, private communication. 14 Clark and Pickett, J. Amer. Chem. SOC., 1930, 52, 465. 15 Taube and Bray, J. Amer. Chem. Soc., 1940, 62, 3357. 16 Fortune and Mellon, Ind. Errg. Chem. (Anal.), 1938, 10, 60. 17 Lindstrand, Svensk Kem. Ticlskr., 1944, 56, 282. 18 Evans and Uri, Trans. Faraday Soc., 1949, 45, 224. 19 Gordon, Hart and Walsh, AECU-1534 (U.S. Atomic Encrgy Commission, de- 20 Weiss, Nature, 1950, 165, 728. 21 Collinson and Dainton, Anir. Rev. Physic. Clzem., 1951, 2, 112. 22 Evans, Hush and Uri, Quart. Rev. (in course of publication). 23 Lefort, Compt. reiid., 1951, 233, 1194. 24 ref. (lo), p. 485. 25 ref. (21), pp. 107-114. 26 Lefort, J . Chim. Phys., 1950, 47, 776. 27 Latimer, Oxidation States of the Elements niid their Potentials in Aqueous Solution (Prentice-Hall, New York, 1938). 28Smith, Anal. Chem., 1951, 23, 925. 29 Haissinsky and Lefort, J. Chim. Phys., 195 1 , 48, 429. 30 Evans and Uri, Nature, 1950, 166, 602. 31 Haissinsky, Lefort and Lebail, J. Chirn. Phys., 1951, 48, 208. 32 Lefort and Haissinsky, J. Chim. Phys., 1951, 48, 368. 33 Silverman and Dodson, BNL-82 (U.S. Atomic Energy Commission, declassified 34 McConnell and Davidson, J. Amer. Cllem. Soc., 1950, 72, 5557. 35 Rabinowitch and Stockmayer, J . Ameu. Chem. Soc., 1942, 64, 335. 36 Igarischew and Turkowskaja, 2. physik. Chem. A, 1929, 140, 227. 37 Dewhurst, J . Chem. Physics, 1951, 19, 1329. 38 Kolthoff and Medalia, J. Arner. Chem. Sac., 1949, 71, 3784. 39 Dainton, private communication. classified document). document), p. 54.C . B. AMPHLETT 155 40 Eimer and Medalia, BNL-117 (U.S. Atomic Energy Cdmmission, declassified 41 see also Barb, Baxendale, George and Hargrave, Trans. Faraday Soc., 1951, 47, 608. 42 Hardwick, this Discussion, ref. (14). 43 Sutton, this Discussion. 44 Rigg, Stein and Weiss, Proc. Roy. SOC. A , 1952, 211, 375. document), p. 33.