20134
J. Phys. Chem. 1996, 100, 20134-20137
Water-Induced Disproportionation of Superoxide Ion in Aprotic Solvents Yong Che, Manabu Tsushima, Futoshi Matsumoto, Takeyoshi Okajima, Koichi Tokuda, and Takeo Ohsaka* Department of Electronic Chemistry, Interdisciplinary Graduate School of Science and Engineering, Tokyo Institute of Technology, 4259 Nagatsuta, Midori-ku, Yokohama 226, Japan ReceiVed: August 19, 1996X
The water-induced disproportionation of the electrogenerated superoxide ion (O2-) in acetonitrile, dimethylformamide, and dimethyl sulfoxide media containing various concentrations of water as a Brønsted acid has been examined by UV-vis spectroscopy. Analysis of the kinetics as a function of O2- and water concentrations and of the measurement time demonstrated that the disproportionation of O2- by water in these media obeys a common mechanism: O2- + H2O a HO2• + OH- (k1, k-1) HO2• + O2- f HO2- + O2 (k2) (HO2•, hydroperoxyl radical; HO2-, hydrogen peroxide anion). The solvent dependence of the obtained kinetic parameters of (k1/k-1)k2, k1 and k-1/k2 is discussed in terms of the solvation of O2- and H2O as well as the effective acidities of H2O in different aprotic solvents. In dipolar aprotic solvents superoxide ion (O2-) is quite stable, because disproportionation to give the peroxide dianion (O22-) is highly unfavorable.1,2 However, the addition of acidic substrates (HA), which act as a Brønsted acid, to stable solutions of O2- in aprotic solvents accelerates the disproportionation, depending on the protic strength (acidity) of HA.1,2 Sawyer et al.3 proposed a mechanism involving H atom transfer disproportionation of hydroperoxyl radical (i.e., HO2• + HO2• f H2O2 + O2) for various acidic substrates in acetonitrile (AN) and in dimethylformamide (DMF). Recently, Save´ant et al.4 have elucidated the mechanism of O2- disproportionation in dimethyl sulfoxide (DMSO) in the presence of weak acids (phenol, substituted phenols, and nitromethane) using double potentialstep chronoamperometry. Their quantitative analysis demonstrates that the HO2• radical undergoes an electron-transfer reduction by O2- itself (i.e., HO2• + O2- f HO2- + O2) rather than abstracting a hydrogen atom from the solvent (i.e., HO2• + SH(solvent) f H2O2 + S•) or exchanging an H atom with another HO2• radical. The present study has been directed to the spectroscopic characterization of the disproportionation of O2- by water as a proton source in three different aprotic solvents (AN, DMF, and DMSO), especially in order to make clear what the mechanism of the O2- disproportionation in aprotic solvents containing various amounts of water is and to compare it with that in aprotic solvents with dissolved Brønsted acids other than water. There have been only a few papers3,5 concerning the water-induced disproportionation of O2- in aprotic solvents, where the quantitative details of the mechanism have not been examined. Experimental Section Reagents. DMSO, AN, and DMF of spectroscopic quality were obtained from Kanto Chemical Co., Inc. Tetraethylammonium perchlorate (TEAP) of reagent grade (Wako Pure Chemicals Industries, Ltd.) was used as the supporting electrolyte. Glassy carbon (GC) disks and plates (GC-20, Tokai Carbon Co., Ltd.) were used as the working electrodes. Water was purified by passage through a Milli-Q purification train. All the other chemicals were used as received. X
Abstract published in AdVance ACS Abstracts, November 15, 1996.
S0022-3654(96)02552-X CCC: $12.00
Instrumentation and Measuring Procedures. Prior to each spectroscopic measurement, solutions of O2- in AN, DMF, or DMSO were prepared by a galvanostatic reduction of O2 using a standard, two-compartment H-type Pyrex glass cell in which cathodic and anodic chambers were separated from each other by a glass frit (G4). A sodium chloride saturated Ag/AgCl reference electrode was connected to the cathodic chamber via a double salt bridge. A GC plate (area 3.2 cm2) and a spiral platinum wire were used as the working and counter electrodes, respectively. The AN (or DMF, DMSO) solution (10-13 mL) containing 0.1 M TEAP and an appropriate amount of water was put in each chamber, and then molecular oxygen (99.98%) was bubbled into the cathodic chamber to obtain almost a saturated solution. Before entering the chamber, oxygen was passed through a trap containing each solvent. Oxygen gas was bubbled throughout the electrolysis. A galvanostatic electrolysis with a constant current of -1.0 mA was performed using a Polarization Unit PS-07 (Toho Technical Co.) in a closed drybox. The amount of charge passed was ca. 0.6-1 C, and the concentration of O2- was typically in the range 0.5-1 mM. The cyclic voltammogram for the redox reaction of the O2/ O2- couple was obtained before each preparative electrolysis of O2- to know a sketchy outline of its redox behavior in each medium in comparison with the previous studies.6-8 The cyclic voltammetric experiments were carried out with a computercontrolled electrochemical measuring system (Model CS1090, Cypress System, Inc.),7 where the working electrode was a GC disk with diameter of 1 mm and the reference and counter electrodes were the same as those used in the above-mentioned preparative electrolysis. In the UV-vis absorption measurements, the O2-saturated O2- solution, prepared as mentioned above, was charged into the quartz cell with the light-path length of 1 cm immediately after the preparative electrolysis, and then the spectrum of O2and its time-course were recorded against a reference (O2saturated solution containing 0.1 M TEAP) at laboratory temperature (23 ( 2 °C) with a Hitachi U-3300 spectrophotometer. The spectra could be recorded up to ca. 190, 260, and 270 nm in AN, DMSO, and DMF, respectively. The kinetic parameters (k1/k-1)k2, k1, and (k-1/k2) for the O2- disproportionation were evaluated in the usual manner from the observed decay of absorbance for O2- at 300 nm using the method of least squares. The production of hydrogen peroxide as a result © 1996 American Chemical Society
Water-Induced Disproportionation of Superoxide Ion
J. Phys. Chem., Vol. 100, No. 51, 1996 20135 steady-state kinetics for hydroperoxyl radical (HO2•), the rate equation reads
-(d[O2-]/dt) ) 2k1k2[O2-]2[H2O]/(k-1[OH-] + k2[O2-]) (3) In addition, we obtain the following equation from mass balance:
-(d[O2-]/dt) ) 2(d[OH-]/dt)
(4)
Here it is implicitly assumed that OH- is generated only via reaction 1 and not by other probable reactions.25 Integrating,
[OH-] ) [OH-]0 + (1/2)([O2-]0 - [O2-])
(5)
where [O2-]0 and [OH-]0 represent the initial concentrations of O2- and OH-, respectively. Substituting eq 5 into eq 3, we obtain Figure 1. Absorption spectra at (1) 0, (2) 2, and (3) 4 h after the preparative electrolysis of O2-saturated 0.1 M TEAP/AN solution containing 41.1 mM H2O. The inset shows a time-course of absorbance at 300 nm during the first 2 h period after the preparative electrolysis (i.e., between spectra 1 and 2).
of the O2- disproportionation was spectroscopically confirmed by measuring the absorption of the yellow solution of the complex formed by H2O2 with titanium(IV) chloride.9 The content of water in each electrolysis solution was monitored with an AQ-7 aquacounter (Hiranuma Co., Ltd.). Results Figure 1 shows typical UV-vis spectra of O2-, which was electrogenerated in a 0.1 M TEAP/AN solution containing 41.1 mM H2O. From this figure it is obvious that the maximum absorption wavelength λmax is 250 nm and is almost comparable to those (249-255 nm) previously reported for O2- in AN2,10-13 and that the absorbance gradually decreases with time owing to the water-induced disproportionation (vide infra). The observed optical absorption is ascribed to 2πg f 2πu electric transition of O2-.14-16 In this case, contrary to the absorption of O2-, the absorbance below ca. 200 nm increased with time. This absorption was ascertained to be due to hydrogen peroxide,13,17,18 which is the product of the O2- disproportionation, on the basis of the similarity of this absorption to that of the AN solution containing hydrogen peroxide as well as the specific reaction of hydrogen peroxide with Ti4+ (see the Experimental Section).9 In further experiments to elucidate the mechanism of the disproportionation of O2-, a time-course of the absorbance at 300 nm was registered to avoid the overlapping of the absorption of O2- with that of hydrogen peroxide and in the case of the experiments in DMF and DMSO to prevent interference from the absorption of the solvents.19 As is known, protonation initiates the disproportionation of O2- in solutions.1,2,21,22 A well-known mechanism of this reaction in aprotic solvents containing acidic substrates (HA) as Brønsted acids can be represented by4,5,23,24 k1
O2- + HA {\ } HO2• + Ak -1
•
-
k2
(1)
-(d[O2-]/dt) ) 4k1k2 [O2-]2[H2O]/{2k-1[OH-]0 + k-1[O2-]0 + (2k2 - k-1)[O2-]} (6) Under the condition that the concentration of H2O does not change appreciably as O2- reacts (i.e., H2O is in large excess over O2-), we can alternatively write
-(d[O2-]/dt) ) a[O2-]2/(b + c[O2-])
(7)
in which a ) 4k1k2[H2O], b ) k-1([O2-]0 + 2[OH-]0), and c ) 2k2 - k-1. Integrating eq 7, the following equation representing the change of [O2-] with time is given:
b[O2-]-1 - c ln([O2-]) - d ) at
(8)
with d ) b[O2-]0-1 - c ln([O2-]0). By rewriting [O2-] in eq 8 by its absorbance A()l[O2-], ) molar absorption coefficient, l ) light-path length), we obtain
{(b′/a)A-1} - {(c/a)ln A} - (d′/a) ) t
(9)
where b′ ) lb, d′ ) lb(A0)-1 - c ln(A0), and A0 ) l[O2-]0. Figure 2A,B shows typical examples of the time-course of absorbance at 300 nm in AN and DMSO, respectively. We can see that in each case the observed decay of absorbance is well fitted, over the entire time-scale of measurement, to the A-t curve calculated based on eq 9 with appropriate values for b′/a, c/a, and d′/a and at the same time that the disproportionation of O2- in the present case obeys the mechanism expressed by eqs 1 and 2.27 In this study, [O2-] was in the range of ca. 0.5-1 mM and [H2O] was ca. 0.17-1 M. On the basis of the effective pKa′ values26 of H2O in the solvents used, it is assumed that [O2-]0 . [OH-]0 and so b′ ) l{k-1([O2-]0 + 2[OH-]0)} ≈ lk-1[O2-]0 ) k-1A0. Note that it is not necessary to previously know [O2-]0 in the present curve-fitting analysis of the observed decay of absorbance to obtain the kinetic parameters; that is, the value of [O2-]0 can be estimated from the extrapolation of the curve-fitted absorbance-time curve to t ) 0. The finally obtained kinetic parameters are summarized in Table 1. Discussion
-
HO2 + O2 98 HO2 + O2
(2)
Assuming that a water-induced disproportionation of O2- may also obey this mechanism (i.e., HA ) H2O) and considering
Comparison with Previous Results. The present kinetic analysis concerning the disproportionation of O2- by water in AN, DMF, and DMSO demonstrates that this reaction obeys a mechanism by which the disproportionation of O2- in aprotic
20136 J. Phys. Chem., Vol. 100, No. 51, 1996
Che et al.
Figure 2. Typical decay curves of absorbance for O2- at 300 nm in (A) 0.1 M TEAP/AN solution containing 41.1 mM H2O and (B) 0.1 M TEAP/DMSO solution containing 55.8 mM H2O. On the basis of eq 9 with b′/a ) 1.44 × 102 s, c/a ) 3.67 × 103 s, and d′/a ) 6.32 × 103 s in (A) and b′/a ) 6.98 × 104 s, c/a ) 8.74 × 103 s, and d′/a ) 1.61 × 104 s in (B), each solid line was obtained from a nonlinear best-fit analysis of all the data (correlation coefficient, 0.98 (A), 0.99 (B); A0, 0.215 (A), 0.163 (B)).
TABLE 1: Kinetic Parameters for the Disproportionation of O2- by Water in AN, DMF, and DMSO solventa AN DMF DMSO
a
[H2O] (M)
103K1k2 (M-1 s-1)
103k1 (M-1 s-1)
k-1/k2
0.0411 ( 0.0021 0.209 ( 0.009 1.00 ( 0.03 0.0182 ( 0.0009 0.0523 ( 0.0024 0.0175 ( 0.0011 0.0558 ( 0.0032 0.128 ( 0.005
8.9 ( 0.4 9.1 ( 0.4 9.4 ( 0.5 2.2 ( 0.1 2.6 ( 0.1 0.066 ( 0.003 0.071 ( 0.004 0.070 ( 0.003
2.8 ( 0.2 2.9 ( 0.2 3.0 ( 0.2 1.9 ( 0.1 2.2 ( 0.1 0.27 ( 0.02 0.29 ( 0.02 0.21 ( 0.01
0.31 ( 0.02 0.32 ( 0.02 0.32 ( 0.02 0.86 ( 0.05 0.85 ( 0.04 4.1 ( 0.3 4.1 ( 0.3 3.0 ( 0.2
Containing 0.1 M TEAP. b K1 ) k1/k-1.
solvents containing various acidic substrates other than water has been recently unambiguously explained by Save´ant et al.;4 that is, the mechanism of disproportionation involves a homogeneous electron transfer between O2- and HO2• and not an H atom exchange between two HO2• radicals. The former amounts to an electron-transfer disproportionation, because HO2• and O2have formally the same oxidation state. The latter mechanism was proposed on the basis of driving force arguments by Sawyer et al.3,28-30 The apparent rate constants K1k2 ()(k1/k-1)k2) of the disproportionation largely depend on the solvent, but appear to be independent of the concentration of water. Their average values are 9.1 × 10-3, 2.4 × 10-3, and 6.8 × 10-5 M-1 s-1 in AN, DMF, and DMSO, respectively. Similarly, the rate constants k1 for the protonation of O2- by water depend on the solvent: the average values are 2.9 × 10-3 (AN), 2.1 × 10-3 (DMF), and 2.8 × 10-4 (DMSO) M-1 s-1. A value of 1.0 × 10-3 M-1 s-1 has been previously obtained for the disproportionation of O2- by water in DMF.3 The values of 3.45 × 10-3 and 4.78 × 10-4 M-1 s-1 have been also reported for the protonation of O2- by water in AN and DMF, respectively.5 A kinetic parameter of k-1/k2 is a measure of the probability for the electron-transfer disproportionation once the protonation of O2- occurred. When k-1/k2 is large, the backward reaction of eq 1 is faster than the electron-transfer disproportionation (eq 2) and the reaction of eq 1 acts as a preequilibrium vis-a`vis the reaction of eq 2, which is the rate-determining step. Conversely, when k-1/k2 is small, the reaction of eq 1 is the rate-determining step. According to the definition for kinetic mechanisms by Save´ant et al.,4 the former limiting behavior is denoted DISP2 and the latter DISP1. The estimated values of k-1/k2 are 0.32, 0.86, and 4.1 in AN, DMF, and DMSO, respectively. Thus, the disproportionation of O2- by water in DMSO is considered as a DISP2 behavior as that of O2- by
phenols in DMSO.4 In addition, we can see that the kinetic regime in AN is closer to the DISP1 case than in DMF and that the kinetic control is no longer by the electron-transfer disproportionation step alone and jointly involves the protonation of O2- radical. A DISP1 behavior has been proposed for the disproportionation of O2- by nitromethane in DMSO.4 Structure of Solvated Water in Different Organic Solvents. The solvent effect on the rate of protonation of O2- will be considered from a viewpoint of solvation of O2- and H2O. The formal redox potentials E°′(O2/O2-) of the O2/O2- couple are -0.54 (DMSO), -0.62 (DMF), and -0.63 (AN) V vs SHE.1,2 If we assume that the E°′(O2/O2-) values are affected primarily by the degree of solvation of O2- (that is, the solvation energy for O2 is assumed to be small and about the same for different solvents), then the relative solvation energies for O2- are, though the difference would be small, of the following order: DMSO > DMF > AN. On the other hand, it is well-known that the solvation of H2O largely depends on the solvent.31-37 Under the present condition that the mole fraction of H2O is in the range of ca. 1 × 10-3 to 1 × 10-2 the enthalpy of mixing of DMSO (and DMF) with H2O is exothermic, while that of AN with H2O is endothermic.37 It is likely that the H2O molecule significantly interacts with DMSO (DMF) molecules via a hydrogen-bond formation between the oxygen atoms (with unshared electron pairs) of the sulfoxide (and amide) moieties in DMSO (and DMF) and H2O, resulting in a rather rigid polymeric structure with an increased viscosity.31,37 Thus, the water is less able to interact with O2- radical anion; that is, “deactivation” of the water occurs, as it is firmly fixed in the polymeric network. That is, such an extremely strong and unique interaction of water with DMSO and DMF would make water protons less available as donors. In addition, it should be noted that mixing of DMSO with H2O is more exothermic than that of DMF with H2O.37 Hence, the stronger interaction of water with DMSO than with DMF results in its lower availability for protonation in the DMSO-H2O mixture than in the DMF-H2O mixture. In the AN-H2O mixture, on the other hand, the nitrile group of AN is not very reactive toward water, and comparatively little deactivation of the water molecules occurs.31,37 In this case, the protonation of O2- appears to proceed relatively easily in AN. We thus believe that such an idea of different solvation of O2- and H2O in AN, DMF, and DMSO reasonably can explain the observed solvent dependence of the protonation rate of O2-: AN > DMF > DMSO. A similar idea has previously been
Water-Induced Disproportionation of Superoxide Ion proposed for the protonation of O2- by water or ethanol in DMF and AN.5 Acidity and Basicity in Different Organic Solvents. A correlationship between the rates for the protonation of O2- and the acidities of various proton donors in a given aprotic solvent has been previously discussed.3,5,22,38 Nanni et al.22 found that the protonation rate of O2- in DMF increases with decreasing the effective pKa′ values of the proton donors (e.g., in going from 1-butanol (k1 ) 1.5 × 10-3 M-1 s-1, pKa′ ) 33.3) to phenol (k1 ) 1.6 × 103 M-1 s-1, pKa′ ) 19.8). On the basis of the data of Save´ant et al.,4 we can also see that k1 tends to increase with decreasing pKa′ of phenol and substituted phenols in DMSO. On the contrary, James and Brown38 reported that there is no relationship between the rate constants for the protonation of O2- and the pKa′ values of proton donors (acetic acid, benzoic acid, phenylacetic acid, phenol, 2-chlorophenol, and 2,6-dimethylphenol) in DMF. From the data obtained in DMF containing H2O or ethanol, Afanas’ev and Kuprianova5 concluded that it is impossible to decide whether or not the protonation rates of O2- depend on the acidities of H2O and ethanol. The similar discussion about k1-pKa′ correlation cannot be straightforwardly applied to the data obtained in different solvents, because the pKa′ values for a given acid in different solvents do not rigorously provide a valid measure of relative proton acidity owing to the assumptions and approximations inevitably involved in their experimental estimation.39,40 By taking into account this point, Barrette et al.26 have estimated the effective acidities (pKa′) for Brønsted acids in aprotic solvents. According to them, the error in the pKa′ values due to junction potential effects in AN, DMF, and DMSO is at most (0.5 pKa′, and the pKa′ values for H2O are 36.7 in DMSO, 34.7 in DMF, and 30.4 in AN. It therefore appears that the k1 values obtained are correlated with the acidity of water in each solvent; that is, k1 increases with decreasing pKa′ for H2O. Acknowledgment. The present work was financially supported by Grants-in-Aid for Scientific Research in Priority Areas of “New Development of Organic Electrochemistry” (No. 236/ 0751221) and “New Polymers and Their Nano-Organized Systems” (No. 277/08246219) from the Ministry of Education, Science, Sports and Culture, Japan, the Nissan Science Foundation, and the Terumo Life Science Foundation. References and Notes (1) Sawyer, D. T. Oxygen Chemistry; Oxford University Press: New York, 1991. (2) Afanas’ev, I. B. Superoxide Ion: Chemistry and Biological Implications; CRC Press: Boca Raton, FL, 1989; Vol. 1. (3) Chin, D. H.; Chiericato, G., Jr.; Nanni, E. J., Jr.; Sawyer, D. T. J. Am. Chem. Soc. 1982, 104, 1296. (4) Andrieux, C. P.; Hapiot, P.; Save´ant, J. M. J. Am. Chem. Soc. 1987, 109, 3768. (5) Afanas’ev, I. B.; Kuprianova, N. S. Int. J. Chem. Kinet. 1983, 15, 1057. (6) Ohsaka, T.; Tsushima, M.; Tokuda, K. Bioelectrochem. Bioenerg. 1993, 31, 289.
J. Phys. Chem., Vol. 100, No. 51, 1996 20137 (7) Tsushima, M.; Tokuda, K.; Ohsaka, T. Anal. Chem. 1994, 66, 4551. (8) Ohsaka, T.; Tsushima, M.; Okajima, T.; Tokuda, K. Denki Kagaku (J. Electrochem. Soc. Jpn.) 1994, 62, 1300. (9) Charlot, G. Chimie Analytique QuantitatiVe; Masson: Paris, 1974; Vol. 2, pp 474, 545. (10) Afanas’ev, I. B.; Kuprianova, N. S.; Polozova, N. I. Int. J. Chem. Kinet. 1983, 15, 1045. (11) Sawyer, D. T.; Calderwood, T. S.; Yamaguchi, K.; Angelis, C. T. Inorg. Chem. 1983, 22, 2577. (12) Fee, J. A.; Hildenbrand, P. G. FEBS Lett. 1974, 39, 79. (13) Ozawa, T.; Hanaki, A.; Yamamoto, H. FEBS Lett. 1977, 74, 99. (14) Krauss, M.; Newmann, D.; Wahal, A. C.; Das, G.; Zemke, W. Phys. ReV. A 1973, 7, 69. (15) Lukin, L. V.; Yakovlev, B. S. Chem. Phys. Lett. 1976, 42, 307. (16) Sawada, U.; Hoyroyed, R. A. J. Chem. Phys. 1979, 70, 3586. (17) Ozawa, T.; Hanaki, A. Chem. Pharm. Bull. 1981, 29, 926. (18) Afanas’ev, I. B.; Kuprianova, N. S.; Letuchaia, A. V. In Oxygen Radicals in Chemistry and Biology; Bors, W., Saran, M., Tait, D., Eds.; Walter de Gruyter: Berlin, 1984; p 17. (19) It has been recently shown that the attack by HO2- of acetonitrile and the subsequent hydrolysis of the adducts result in the products with optical absorption at ca. 254 and 290 nm.20 No direct observation of such an absorption was made under the present experimental conditions. (20) Yamaguchi, K.; Caldewood, T. S.; Sawyer, D. T. Inorg. Chem. 1986, 25, 1289. (21) Sawyer, D. T.; Vallentine, J. S. Acc. Chem. Res. 1981, 14, 393. (22) Nanni, E. J., Jr.; Stallings, M. D.; Sawyer, D. T. J. Am. Chem. Soc. 1980, 102, 4481. (23) Frimer, A. A.; Marks, V.; Gilinsky-Sharon, P. Free Rad. Res. Commun. 1991, 12-13, 93. (24) Frimer, A. A. In The Chemistry of Peroxides; Patai, S., Ed.; Wiley: Chichester, 1983; pp 429-461. (25) For example: (i) H2O- + H2O U OH- + H2O2, (ii) H2O U H+ + OH-, and (iii) B- + H2O U HB + OH- (HB, impurity), pKa values of OH- and HO2- in water (14.0 and 11.8) may suggest that the equilibrium of (i) largely lies to the left also in aprotic solvents such as AN, DMF, and DMSO. The effective acidities (pKa′)26 of H2O in aprotic solvents are very low compared with that (14.0) in water, e.g., pKa′ ) 30.4 in AN, 34.7 in DMF, and 36.7 in DMSO. Thus, the dissociation (ii) of H2O in these aprotic solvents is negligible. When a base stronger than H2O is present as an impurity, OH- might be generated on the basis of reaction (iii). The presence of such an impurity is also negligible in the present experiment. (26) Barrette, W. C., Jr.; Johnson, H. W., Jr.; Sawyer, D. T. Anal. Chem. 1984, 56, 189. (27) The observed decay of absorbance of O2- could not be fitted to the absorbance-time curve calculated on the basis of an H atom transfer disproportionation mechanism.2,28,29 (28) Roberts, J. L.; Sawyer, D. T. Isr. J. Chem. 1983, 23, 430. (29) Cofre, P.; Sawyer, D. T. Anal. Chem. 1986, 58, 1057. (30) HO2- is a much stronger base than O2- (pKa values for H2O2 and HO2• are 11.8 and 4.9, respectively, in H2O3,4). Thus, it seems also likely that reaction 2 is followed by the reaction HO2• + HO2- f H2O2 + O2-; that is, the overall reaction is 2HO2• + O2- f H2O2 + O2- + O2. (31) Jezorek, J. R.; Mark, H. B., Jr. J. Phys. Chem. 1970, 74, 1627. (32) Vierk, A. Z. Anorg. Chem. 1950, 261, 285. (33) Tommila, E.; Murto, M. Acta Chem. Scand. 1963, 17, 1947. (34) Cowie, J. M. G.; Toporowski, P. M. Can. J. Chem. 1961, 39, 2240. (35) Ting, S. F.; Wang, S. M.; Li, N. C. Can. J. Chem. 1967, 45, 425. (36) Ponner, O. D.; Choi, Y. S. J. Phys. Chem. 1974, 78, 1723. (37) Maucus, Y. Ion SolVation; John Wiley & Sons Ltd.: Chichester, 1985; Chapter 7. (38) James, H. J.; Broman, R. F. J. Phys. Chem. 1971, 75, 4019. (39) Popvych, O.; Tomkins, R. P. T. Nonaqueous Solution Chemistry; Wiley: New York, 1981; p 165. (40) Popvych, O. Anal. Chem. 1974, 46, 2009.
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