At pH > 4, the oxidation is inhibited by organics, suggesting a free radical mechanism. One proposed mechanism, which originates in the work of Backstrom (1934), is shown in Table 8.7. The inhibition occurs when the organics react with the sulfate radical ion, S04~. This inhibition has also been seen in laboratory experiments using fogwater collected in Diibendorf, Switzerland, where oxidation rates for S(IV) were less than expected based on the kinetics of the iron-catalyzed oxidation (Kotronarou and Sigg, 1993).

The Mn2+-catalyzed oxidation of S02 is also complex in both kinetics and mechanism. At pH 2, for example, the reaction is second order in Mn2+ and zero order in S(IV) (i.e., is independent of the dissolved sulfur concentration) when [S(IV)] > 10~4 mol L~ 1 but first order in both Mn2+ and S(IV) at low concen-

TABLE 8.7 Proposed Mechanism for the Catalyzed Oxidation of S(IV) in Aqueous Solutions of pH Range 4-7"

S04~ -> products kh kh = 4.5 X 109 1 s~ 1 kc = 1.5 x 109 M~1 s~1 kd = 1.3 X 107 M~ 1 s~ 1 kt. = 2 x 109 M~ 1 s~ 1 kf= 9.9 x 10s M~ ' s~ ' kg = see Neta et al. (1988) and Wine et al. (1989) kh = 410 s"1

" Backstrom (1934), Huie and Neta (1987), and Martin (1994). Rate constants from Huie and Neta (1987), Neta et al (1988), Tang et al, (1988), and Wine et al. (1989).

trations, [S(IV)] < 10 6 mol L Furthermore, the rate decreases with ionic strength at all S(IV) concentrations. Finally, there is a synergistic effect in the presence of both Mn2+ and Fe3+ (Martin, 1994).

In summary, the uncatalyzed oxidation of S(IV) occurs in aqueous solution but is very slow. However, given the ubiquitous occurrence of Fe3+ and Mn2+ (see Chapter 9), the uncatalyzed oxidation is likely irrelevant to atmospheric solutions. The catalyzed oxidations are complex in both kinetics and mechanism. We shall defer a comparison of their importance until other oxidation mechanisms are discussed. However, we shall see that the catalyzed oxidations are likely to contribute significantly to S(IV) oxidation in solution only at pH values near neutral, i.e., in the range of ~6-7. As the oxidation occurs and acid forms, the pH falls. The rapid falloff in the rate of the catalyzed oxidation with increasing [H + ] then results in a rapid quenching of this path, as expected from Fig. 8.9a.

As discussed in Chapter 7, there is some evidence that freezing of aqueous solutions containing nitrite accelerates its oxidation to nitrate (Takenaka et al., 1992, 1996). A similar phenomenon has been reported in cloud chamber studies, where sulfide was observed to be oxidized to sulfate during ice crystal formation from expansion of droplets containing ionic salts (Finnegan et al., 1991; Finnegan and Pitter, 1991; Gross, 1991). Possible mechanisms are discussed by Finnegan and Pitter (1997).

d. Oxidation by 03

While the Henry's law constant for ozone is fairly small (Table 8.1), there is sufficient ozone present in the troposphere globally to dissolve in clouds and fogs, hence presenting the potential for it to act as a S(IV) oxidant. Kinetic and mechanistic studies for the 03-S(IV) reaction in aqueous solutions have been reviewed and evaluated by Hoffmann (1986), who shows that it can be treated in terms of individual reactions of the various forms of S(IV) in solution. That is, S02 • H20, HSO^", and SO2- each react with 03 by unique mechanisms and with unique rate constants, although in all cases the reactions can be considered to be a nucleophilic attack by the sulfur species on 03.

Figure 8.14 shows the proposed mechanisms of reaction for each species. The overall rate of the S(IV) oxidation can then be represented by

where a0, a,, and a2 are the fractions of the total S(IV) in the form of S02 • H20, HS03", and S032", respectively. That is, a(l = [S02 • H2Oj/[S(IV)] etc., where S(IV) is the sum of {S02 • H20 + HS03" +


03 - Aquated S02 Mechanism

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