J. CHEM. SOC. FARADAY TRANS., 1994, 90(21), 3309-3313 Reaction of Peroxornonosulfate Radical with Manganese(![) in Acidic Aqueous Solution A Pulse Radiolysis Study J. Berglund and L. 1. Elding Inorganic Chemistry 1, Chemical Center, P.O. Box 124,22100 Lund, Sweden G. V. Buxton, S. McGowan and G. A. Salmon" Cookridge Radiation Research Centre, The University of Leeds, Cookridge Hospital, Leeds, UK LS16 6QB The reaction between the SO5-radical and Mn" has been proposed to be the most important process for regeneration of Mn"' in the Mn'"/''-catalysed autoxidation of S" in acidic aqueous solution. In the present study, the second-order rate constant for this reaction has been determined at pH 3.0 and 10 mmol dm-, ionic strength by use of pulse radiolysis. The study was performed in the presence of excess S".Under these conditions Mn" is distributed among the complexes Mn2'(aq), [Mn(HSO,)]+ and [Mn(S03)Mn12+. The rate of reaction decreases as a function of increasing [Mn' which is rationalized qualitatively by a mechanism involving three parallel reactions between SO5-and the Mn" complexes, with rate constantsk,, , k,, and k,, , respectively. H+Mn2++ SO5--Mn3++ HSO,-(16) H+ [Mn(HSO,)]+ + SO,--[Mn(HSO3)l2+ + HSO,-(17) H+[Mn(S0,)Mnl2+ + SO,--[Mn(S03)Mn13' + HS0,-(18) For increasing total concentrations of Mn", formation of the sulfito-bridged complex is favoured which implies that k,, < k,, , k,, . Values of the second-order rate constant in the range 2 x lo8-2 x 10,' dm3 mol-' s-' have been determined, depending on which Mn" species is predominant.Subsequent slow processes are observed following the formation of Mn"'. These reactions have been attributed to the disproportionation of Mn"' and reactions between the Mn"' species and excess S". The implications of the present results for the Mn'l'/l' catalysed autoxidation of S'' are discussed. Recent results indicate that aqueous phase oxidation of mechanism of the reaction between hexaaquamanganese(I1) sulfur(rv)(SO,. nH,O, HSO,-, SO,,-) by molecular oxygen, and the peroxomonosulfate radical with pulse radiolysis in catalysed by metal ions, uiz. MnIrl/I1, Fel*I/I1, Co"'/" and CulI1/I1, order to determine the rate constant for this reaction and takes place by a common free-radical chain mechanism. gain a better understanding of the manganese-catalysed The chain is initiated by reaction between the trivalent metal autoxidation of SIV.ion and SIV according to reaction (1) and is propagated by SO,-, SO5-and SO4-radicals. Experimental M"' + HSO,--M" + SO,-+ H+; Chemicals and Solutions (M = Mn, Fe, Co, Cu) (1) A stock solution of 2.72 mmol dm-, manganese@) perchlor- ate was prepared by dissolving an accurately weighed The reduced metal ion is re-oxidized by the strong oxysulfur amount of Mn(ClO,), -H20 (Johnson Matthey GmbH) in radicals SO5-and SO4-as well as by hydrogen per-100 ml water. The pH of the secondary manganese@) solu- oxornonosulfate, HS05-, which is also generated in the tions was adjusted to pH 3.0 by use of perchloric acid (BDH,chain, reactions (2)-(4).60%)and the ionic strength was kept at 10 mmol dmP3 using sodium perchlorate (BDH, 99.9%) as supporting electrolyte. H+M" + SO,--M"' + HS05-(2) The solutions were saturated with oxygen by use of the stan- dard syringe-bubbling technique as described previo~sly.~ M" + SO4--M"' + MI' + HS05--M"' + SO4-+ OH-(3) An oxygen-free stock solution of ca. 20 mmol dm-, (4) sulfur(1v) was freshly prepared before each set of experiments by dissolving ca. 0.26 g Na,SO, (Merck p.a.) in 100 ml 20 Consequently, these reactions are very important in order to mmol dm- perchloric acid, which was continuously flushed close the catalytic cycle and to rationalize the complicated with argon. The concentration was checked by use of stand-processes governing metal ion catalysed autoxidation of S".ard ion chromatography and found to be constant to within However, the kinetics and mechanisms for most of these reac- 97% of the prepared value. The pH and ionic strengths of the tions have not been studied. secondary sulfur(1v) solutions were adjusted to 3.0 and 10 Recently, it has been proposed that reaction (2) is mmol dm-,, respectively, and they were saturated with responsible for the regeneration of Mn"' in the manganese- nitrous oxide using the syringe-bubbling technique. catalysed process and that reactions (3) and (4)can be All solutions were prepared by use of deionized (Millipore, neglected, provided reaction (2) is fast, i.e. k, 9 lo4 dm3 Milli-Q) water and all gases were used as received from the mol-' s-l.' Therefore, we have studied the kinetics and suppliers without additional purification.J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Apparatus and Measurements Pulse radiolysis experiments were carried out using a pulsed electron beam generated by a 3 MV Van de Graaff acceler- ator. Peroxomonosulfate radicals were produced by irradiat- ing, with 17-41 Gy, 0.2 ps pulses, a solution containing Mn" and SIV,half saturated with N,O and oxygen, and contained in a triple-pass cell of optical length 6.8 cm,see below. In order to minimize thermal oxidation of SIV before irradiation, equal volumes of separate solutions of N,O-saturated S" and 0,-saturated Mn" were mixed in a rapid-mix apparatus less than a second before the pulse was delivered.Thus, in all the experiments, the irradiated solutions were half-saturated with both N,O and O,, i.e. [N,O] = 0.0125 mol dm-3 and [O,] = 675 pmol dm-3. A xenon arc lamp, operated in either continuous or flashed mode, was used as the analysing light source. The dose per pulse was measured using the standard thiocyanate dosimeter (lo-, mol dm- 3, saturated ~with oxygen, for which GE= 2.48 ~x ~ m2 J-'.' The experiments were carried out at ambient temperature, i.e. 20 & 1 "C. Kinetic data were evaluated using either the TREAT program,6 developed to analyse simple pulse radiolysis traces, or the FACSIMILE kinetic modelling ~rogram.~ Generation of SOs-The radiolysis of an aqueous solution results in the gener- ation of hydrated electrons, eaq-, hydroxyl radicals, a small yield of hydrogen atoms and small amounts of hydrogen per- oxide and hydrogen, reaction (5).In the presence of N,O, the hydrated electrons are converted to OH radicals by reactions (6) and (7) with k, = 9.1 x lo9 dm3 mol-'s-' and k, = 9.4 x 107 s-,.~ ksN,O + eaq--N, + 0-0-+ H,O OH + OH-(7) The reaction of OH with Mn" is a hundred times slower than that of OH with hydrogen sulfite and conditions were chosen such that 99% of the OH radicals reacted with HS03-, gen- erating sulfite radicals according to reaction (8) with k, = 4.5 x 10' dm3 mol-'~-~ ksHS03-+ OH -SO3-+ H20 (8) Sulfite radicals react rapidly with dissolved molecular oxygen in a subsequent step forming peroxomonosulfate radicals, reaction (9), with kg = 2.5 x lo9dm3 mol-I s-l.' k9SO3-+ 0, -SO5-(9) SO5-radicals are known to be involved in the chain propa- gation steps of the free-radical chain oxidation of SIV by molecular oxygen, the mechanism of which has been dis- cussed in detail by McE1roy.l' The relevant reactions at pH 3.0 are reactions (10) and (ll), the rate constants of which have been determined to be:" k,, = (1.0 0.1) x lo3 and k,, = (1.1 & 0.1) x lo4dm3 mol-' s-'.An earlier estimate of these rate constants gave only the upper limit, i.e. (klo+ kll) < 3 x lo5dm3 mol-' s-l.', SO5-+ HS0,-+ HS0,-+ SO3-(10) SO5-+ HS03--+ SO4-+ HS04-(1 1) These values indicate that these reactions are unimportant under the conditions and timescales used in the present experiments and that reactions of SO4-may, therefore, be neglected. In addition to the small yield of H generated in the primary radiolytic event, reaction (5), a further small yield is formed by reaction (12) in competition with reaction (6).e-(aq) + H+ 4H (12) At pH 3.0 the total yield of H is calculated to be 21% of the yield of radicals and in the presence of 0, at the concentra- tion used in this study H will be converted to HOz with t1,, x 50 ns. HO, reacts with Mn2+(aq) with a rate constant of 6.0 x lo6 dm3 mol-' s-l.13 Thus at the highest concentra- tion of Mn2+ used in this study (500 pmol dm-3) this reac- tion is 100 times slower than reaction (2). Results and Discussion Preliminary Observations Fig.1 shows a typical kinetic trace at 470 nm after irradiation of Mn"/SIV solution. The rapid increase in absorbance fol- lowed by a somewhat slower decrease is probably due to the formation of 0,-and its protonated form HO, generated in reactions (13) and (14), where reaction (13) is in competition with reaction (Is)? 0-+ 0, -,0,- (13) 0,-+ H+(H,O)eHO,(+OH-) (14) 0-+ H+(H,O) -+ OH(+OH-) (15) We have observed similar long-lived absorptions at both pH 6.0 and 3.0 in pulse-irradiated water containing both N,O and O,, but they are not present in the absence of 0,.14 The rate constants for reactions (12) and (14) indicate that the yield of 0,-is expected to be only 2.9% of that of OH, but 0,-absorbs strongly at 470 nm.1s.16 Further evidence that 0-can react with solutes in systems at low pH is provided by the work of Zehavi and Rabani.l7 time Fig. 1 Typical kinetic trace at 470 nm obtained by pulse radiolysis of an aqueous solution of Mn" and S'", half-saturated with molecular oxygen and nitrous oxide. Conditions: [Mn"] = 100 pmol dm-3, [S'"] = 800 pmol dmP3, pH 3.0, I = 10 mmol dm-3, dose = 17 Gy. t Note added in proof: In recent experiments we have shown that protonation of 0,-at pH 3.0 is too rapid to account for the 'spike' in Fig. 1 which, at present, we attribute to transient absorption induced in the cell. J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 200 ps H IIIIIIII time Fig. 2 Decay of the Mn"' product at 440 nm.Conditions: [Mn"] = 100 pmol dm-3, [StV] = 200 pmol dm-3, pH 3.0, I = 10 mmol dm-3, dose = 30 Gy. The first-order increase in absorbance to a plateau value is attributed to formation of Mn"'. Under certain experimental conditions, see below, a slow decrease in absorbance could be observed following the build-up of Mn'", Fig. 2. Both reactions were studied in the wavelength range 440 to 470 nm at pH 3.0 under the experimental conditions 50 < [Mn"]/pmol dmw3< 500, 200 < [SIV]/pmol dm-j < 800 and I = 10 mmol dm-3. Doses in the range 17-41 Gy were used, corresponding to an initial concentration of SO, -in the range 10-25 pmol dm -'. Spectra Fig. 3 shows spectra recorded at 2, 20 and 170 ys after irra- diation of a mol dm-' Mn" and 2 x mol dm-3 S" solution and presented as GE-values, i.e. the product of 61 I I I I I I 5 -,., 4-1, c 0 I -I 250 350 450 550 wavelength/nm Fig.3 Spectra recorded at (-) 2, (---) 20 and (. . -.) 170 ps after irradiation and, for comparison, (-. the spectrum of SO, -adjust-0-) ed to G(SO,-)= 5.0 x rnol J-'. Conditions: [Mn"] = 100 pmol dm-3, [S"] = 200 pmol dm-j, pH 3.0, I = 10 mmol dmb3, do= = 19 Gy. the radiation chemical yield, G, of the light absorbing species in units of mol J-and E, the species' absorptivity in units of m2 mol-'. After 20 ,us the formation of manganese(II1) is com- plete and the spectrum shows a maximum at ca. 450 nm. The spectrum recorded after 170 ps shows that the absorbance has decreased over the spectral region from 420 to 600 nm owing to the subsequent, slower reaction.Also shown in Fig. 3 is the spectrum of SO, -expressed in terms of GEbased on G(S0,-) = 5.0 x lo-' mol J-I, the expected yield of SO,-in this system. Additional spectra of the manganese(@ product were recorded at the plateau of the kinetic traces, cf: Fig. 1, for solutions containing different [Mn"] and [S"]. All exhibit a maximum at 260 nm and most of the spectra have an addi- tional, but smaller, absorption in the wavelength range 350-600 nm with a maximum in the region 440-470 nm which is characteristic of many Mn"' absorption spectra previously reported." The extent of absorption in the region 440-470 nm is dependent on the [S'"] : [Mn"] ratio.The absorption decreases as the ratio increases, Table 1, and assuming G(Mn"') = 5.0 x lo-' rnol J-', the effective molar absorp- tivity varies from 1.0to 4.2 mz mol-'. This effect may be due to formation of different Mn"' complexes, e.g. hydrolysed manganese(m), [Mn(HS03)lZ and [Mn(S03)Mn13 +,under+ different experimental conditions, see discussion below and ref. 1. These complexes most likely have different absorp- tivities in this wavelength region, but the values of &:Ao lie in the range previously observed for complexes of Mn"'. 19,20 The apparent decrease in absorbance change on the forma- tion of manganese(In), when the [S"] : [Mn"] ratio increases, may also be explained by an increase in the rates of the sub- sequent reactions when the SrV concentration is increased.Kinetics The growth of absorbance due to formation of Mn"' was analysed according to the pseudo-first-order rate law (I). d[Mn"']jdt = k,b,[SO,-] (1) A first-order correction for the subsequent slow decay was applied when necessary. The value of the observed rate con- stant decreases with increasing concentration of Mn" using both 19 and 41 Gy pulses, Table 2 and Fig. 4. There is also a small dependence on the S" concentration. The values of kobs are slightly larger using [S"'] = 800 pmol dm-3 than using [S"] = 200 ymol dm-3, Table 2. kobs is slightly dependent on dose but is independent of wavelength, within experimental error. Expressed as bimolecular rate constants the values range from 2 x lo8 to 2 x 10" dm3 mol-' s-'.It is sug-gested that the variation in these values is due to a change in the nature of the Mn" species with [S"] and [Mn"] (see below). Table 1 GE values at 470 nm for the Mn"' product determined at different concentrations of S"' and Mn" cs'vl /pol dm-3 [Mn"]/pmol dm-3 [S"] : [Mn"] GE470 mz J-' 200 500 0.4 : 1 2.1" 200 200 1:l 1.Sb 800 500 1.6 : 1 1.5" 200 100 2:l 1.4" 800 200 4: 1 1.2b 800 100 8:l 0.90" 5000 500 10: 1 1.o" 5000 100 50 : 1 0.50" " 19 Gy pulses were used at pH 3.0. 41 Gy pulses were used at pH 3.0. I H i I I I I I I 12 r I ma 0 4 0' I I I I I I0 200 400 600 [Mn "]/pmol dm-3 Fig. 4 Observed pseudo-first-order rate constant as a function of excess concentration of Mn".Conditions: see Table 2. [S"] = (a)200 and (b) 800 pmol dm-3, dose = 19 Gy (closed symbols) and 41 Gy (open symbols). The errors are given as the standard deviations of the mean calculated from four to six measurements. Mechanism Recently, it was shown that a complex between Mn" and hydrogen sulfite with stability constant bl = 3 x lo4 dm3 mol-' is formed in the interval 2.4 < pH < 4.0.' At pH 4.0, Table 2 Observed pseudo-first-order rate constants at different con- centrations of S" and Mn" 4,,~105 s-1 CS'"1 [Mn"] /pmol dm-j /pmol dm-3 dose = 19 Gy dose = 41 Gy 200 50 3.4 k 0.5 4.1 & 0.8 200 100 3.0 f 0.3 3.6 k0.7 200 200 2.0 f 0.1 2.1 & 0.4 200 300 1.6 f 0.1 1.5 & 0.1 200 400 1.0 f 0.1 1.3 & 0.1 200 500 1.1 * 0.1 1.5 f0.1 800 50 9.4 f0.1 11.0 f2.5 800 100 9.6 f0.1 8.0 f0.7 800 200 7.7 & 0.3 7.3 f1.0 800 300 6.3 f0.1 5.8 f0.6 8W 400 1.9 & 0.1 2.1 f 0.1 80% 500 2.8 _+ 0.1 The errors are grven as the standard deviations of the mean calcu- lated from four io six measurements.Conditions: pH 3, Z = 0.1 mol dm-3. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 formation of a bridged Mn" complex, [Mn(S03)Mn12+ was also indicated.' Consequently, the distribution of Mn" between various complex species in the present study is gov- erned by the total concentrations of SrV and Mn". The total concentration of Mn" can be written according to eqn. (11) where #Il2 denotes the stability constant of [MnS03Mn12+.[Mnllltotai= [Mn2'1(1 #Ii[HsO3-1 + B12CMn2+ICHS03-I/CH+1) (11) The dependence of kobs on the concentrations of Mn" and STv may be rationalized qualitatively by use of the following mechanism, reactions (16)-( 18). Mn2+ + SO5----+ H+ Mn3+ + HS0,-(16) H+[Mn(HSO,)]+ + SO,--[Mn(HSO3)l2++ HSO,-(17) [Mn(S03)Mn12++ SO,-H+ [Mn(S03)Mn13---+ + + HS05-(18) The observed rate constant of eqn. (I) will be the sum of the contributions from these three parallel reactions with rate constants k16, k17 and k18 and can be written as eqn. (111). kobs = k16[Mn2+] 4-k,,[Mn(HSO,)+] 4-k18[Mn(SO3)Mn2+] (111) As the concentration of Mn" is increased, the ratio [Mn(S03)Mn2 +]/[Mn"],,,,, increases according to eqn.(11) and thus, the rate of reaction decreases provided that k18 < k16, k17. The small increase in the value of kobs observed for increasing S'" concentrations indicates that k, 7 is slightly larger than k16 since a change of the total s" concentration only influences the ratio between [Mn2'] and [Mn(HSO,)+]. Based on a mechanism proposed by Jayson et al. for the reaction between Fe2+ and H0, Cabelli and BielskiI3 have suggested the following mechanism for the reaction between Mn" and the perhydroxyl radical: Mn2+(aq)+ HO, -+ [Mn(OOH)I2+ (19) [Mn(OOH)]2' + Mn2+(aq)+ [Mn(0OH)Mnl4+ (20) [Mn(OOH)I2+-+ Mn3+(aq)+ H0,-(21) [Mn(OOH)MnI4+ + Mn3+(aq)+ products (22) However, in the mechanism proposed by Jayson et al. for the Fe2+/H02 system, Fe2+ is oxidized to Fe3+ by HO,.After electron transfer, the reaction products persist together as an outer-sphere successor, reaction (23). Fe2+ + HO, -+Fe3+, H0,- (23) The successor complex is in equilibrium with a bridged species, [Fe11'H0,Fe"]4+. Both complexes decompose to give Fe3+ and H02-and Fe3+, Fe2+ and HO,-, respec-tively. Thus, reaction (23) is the only redox reaction in the mechanism, contrary to what is claimed by Cabelli and Bielski. Generation of [Mn(00H)I2 + followed by formation of [Mn(00H)MnI4+ according to reactions (19) and (20) does not seem likely since intramolecular electron transfer in [Mn(OOH)I2' is expected to be very fast. Note that in [Mn(OOH)]'+ a radical is coordinated to Mn", while in Fe3+, H0,-it is a molecular ligand that interacts with the oxidized metal centre.The mechanism proposed in the present study rationalizes qualitatively the experimental results without introducing J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 any complex formation between Mn2+ and SO,-. It should be emphasized that a mechanism analogous to the one postu- lated by Cabelli and Bielski cannot explain why the value of the observed rate constant decreases for increasing concen- trations of Mn" unless [Mn(S0,)Mn13 decomposes more + slowly than [Mn(SO,)] +.This implication would be in sharp contrast to the conclusion made by Cabelli and Bielski that [Mn(00H)MnI4+ decomposes considerably faster than [Mn(00H)] ' .+ Subsequent Slow Processes By use of 41 Gy pulses, the decrease in absorbance following the formation of Mn"' was observed using 50 < [Mn"]/pmol dm-' < 500 and [S"] = 200 and 800 pmol dm-3.However, at lower doses, CQ. 19 Gy, this decay was observed only at low Mn" concentrations, i.e. 50 and 100 pmol dm-3. At a constant Mn" concentration (1 x mol dm-3) the decay rate slows down with increasing sulfur(1v) concentration. At low doses, the absorption decays according to simple first- order kinetics while at higher doses, the decays are best described by a second-order process. It was not feasible to rationalize these slow processes quantitatively. However, it should be emphasized that generation of Mn"' according to reactions (16)-( 18) triggers the manganese-catalysed autoxi- dation of S*.The Mn'" species all react further with different rates and form sulfite radicals, reactions (24)-(26). These rad- icals react with molecular oxygen reproducing SO -radicals, reaction (9). Mn3+(aq)+ [Mn(HSO,)]+ 2Mn2+(aq)+ SO3-+ Hf (24) [Mn(HSO3)l2+-+ Mn2+(aq)+ SO3-+ H+ (25) [Mn(S03)Mn13+-+ 2Mn2+(aq)+ SO3-(26) The peroxomonosulfate radical then reoxidizes Mn" to Mn"' and so on. On the short timescales used in the present study, steady-state conditions have not been established. The pro- cesses following the initial formation of Mn"' are therefore complex. In addition to reactions (24)-(26) dispro-portionation of Mn"' according to reaction (27) is also pos- sible. 2Mn3+(aq)+ 2H20eMn02+ 4H+ + Mn'+(aq) (27) Changing the dose and the initial concentrations of Mn" and S'" may therefore favour different pathways for the decay of M n"'.Conclusion The present study clearly demonstrates that Mn" is oxidized to Mn"' by SO,-radicals and that this is a very fast reaction. The second-order rate constants are approximately 2 x lo8-2 x 10" dm3 mol-' s-' depending on which of the Mn" complexes is predominant in solution. This result gives con- clusive support to the suggested mechanism for the manganese-catalysed autoxidation of SIV. The system is complicated owing to formation of sulfito complexes of Mn" and the initiation of the autoxidation of SIV.Further informa- tion about the kinetics and mechanisms for reactions with peroxomonosulfate radicals and Mn" may be provided by a laser flash photolysis study currently planned at Cookridge Radiation Research Centre.This technique allows the rad- icals to be generated in the absence of StVwhich simplifies the observed kinetics. Sulfite radicals are produced by homolytic cleavage of dithionate and the SO,-radicals are formed by subsequent reaction of the sulfite radicals with molecular oxygen according to reaction (9). Financial support from the Commission of the European Communities within the STEP research program (contract STEP-005-C), from the Swedish Natural Science Research Council and from the Royal Physiographic Society of Lund is gratefully acknowledged. Prof. Sture Fronaeus is acknow- ledged for valuable comments.References 1 J. Berglund, S. Fronaeus and L. 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