O

At high pH values (which are not typical of most aqueous solutions in the atmosphere, however), H202 can also react with SOj (Lagrange et al., 1996). Other studies (McArdle and Hoffmann, 1983; Kunen et al.,

1983; Overton, 1985) are also in agreement with the data of Fig. 8.16. It is seen that in contrast to most of the other oxidations discussed earlier, ku decreases as the pH increases at pH > 1.5. This is in the opposite direction to the S(1V) solubility (i.e., is qualitatively described by the situation depicted in Fig. 8.9b). The result is that, unlike the other reactions discussed so far, the overall rate of production of S(Vt) from this reaction is relatively independent of pH over a wide pH range of interest in the atmosphere (see below). Interestingly, Lagrange et al. (1993) report that the reaction is catalyzed by chloride and ammonium ions.

The kinetics of the reaction can be expressed as

-= &[H+][H202][S(IV)], at where k has been measured to be in the range (7.2 + 2.0) X 107 to (9.6 ± 0.5) X 107 L2 mol"2 s"1 (Lee et al., 1986; Lind et al., 1987). Typical H202 levels are up to a maximum of ~ f 0 ppb in the gas phase and more typically ~1 ppb; in cloudwater, concentrations up to ~200 ju,M have been reported (e.g., see Sakugawa et al., f990; Claiborn and Aneja, 1991; and Lee et al., 1993). The ubiquitous occurrence of H202, its large Henry's law constant, its high reactivity, and the pH dependence of the rate constant combine to make H202 the most important oxidant for S(IV) in the troposphere.

This is consistent with the results of a number of field studies. For example, the release of S02 into air entering a cloud base at Great Dunn Fell, England, resulted in the simultaneous increase in sulfate and decrease in H202 in the cloudwater (Clark et al., 1990), and H202 concentrations are often inversely related to S02 (e.g., Sakugawa and Kaplan, 1989).

ft is not clear whether the reaction kinetics will be significantly altered by other species present in the aqueous phase. For example, Lee et al. (1986) showed that the kinetics of the H202-S(tV) reaction in freshly collected precipitation were only 15% below those measured in laboratory pure water. On the other hand, laboratory studies by Lagrange et al. (1993) suggest that the rate constant depends both on the ionic strength and on the nature of the electrolyte and that Fe2+ catalyzes the reaction.

Organic hydroperoxides have also been proposed as potential oxidants of S(tV) in solution (e.g., Graedel and Goldberg, 1983):

ROOH + HSOJ -> HS04" + ROH, (19) kw = 7.2 X 104 Lmor1 s"1.

Both methyl hydroperoxide (CH3OOH) and peracetic acid (CH3C(0)00H) have been shown to oxidize S(1V)

in the aqueous phase (e.g., Lind and Lazrus, 1983; Lind et al., 1987). Lind and co-workers (1987) express the rate law for oxidation of S(IV) by H202, CH3OOH, and CH3C(0)00H in the form

Rate (mol L~1 s"1) = A:[H+][peroxide][S(IV)], with k = 7.2 X 107 L2 mol"2 s"1 for H202, 1.7 X fO7 L2 mol"2 s"1 for CH3OOH, and 3.5 X 107 L2 mol"2 s"1 for CH3C(0)00H. (The rate for CH3C(0)00H was shown to contain an additional term, k [peroxide]-[S(IV)], where k = 610 L moL1 s-1, which becomes important above a pH of ~5.) As seen from the data in Table 8.1, the Henry's law constant for CH3OOH is about three orders of magnitude less than that for H202 and that for CH3C(0)00H is about two orders of magnitude less than for H202. Since their aqueous-phase concentrations are expected to be smaller than that of H202 they are not expected to contribute significantly compared to oxidation by H202. (See, for example, Kelly et al., 1985.)

Zhou and Lee (1992) measured the rate constant for the reaction of hydroxymethyl hydroperoxide, HOCH2OOH, with S(IV) compared to H202 and found *sav,+ .ic>c.i,<x>.i = {2-2 X 107[H + ]} L2 mol"2 s"1 (based on ¿s(iv)+ii,o, = i9-6 x 107[H + ]} L2 mol"2 s"1 determined by Lee et al., 1986). As seen in Table 8.1, this hydroperoxide, which can be formed from the H202-HCH0 reaction or the reaction of H20 with the Criegee intermediate (HCHOO-) (see Chapter 6), has a Henry's law constant that is even larger than that for H202. Because of uncertainties in its concentrations in fogs and clouds, it is not possible to make a firm estimate of its importance in the oxidation of S(tV). However, in one set of measurements made in Georgia, gas-phase HOCH2OOH constituted a large percentage of the total peroxides, with concentrations as high as 5 ppb (Lee et al., 1993). On the other hand, near Grand Canyon, Arizona, most of the total peroxide was measured to be H202 (Tanner and Schorran, 1995). Hence, HOCH2OOH may be important in S(IV) oxidation under some circumstances, depending on the relative amount compared to H202.

f Oxidation by Oxides of Nitrogen

The oxides of nitrogen—NO, NOz, N03, HONO, and HN03—have all been suggested as possible oxidizing agents for dissolved S(1V); however, the reactions of HN03 and NO at atmospheric concentrations are too slow to be significant (Lee and Schwartz, 1983; Martin, 1984; Schwartz, 1984a).

In aqueous solutions, nitrous acid reacts with S(IV) at a reasonable rate with a rate expresssion given by

However, the levels of gaseous HONO observed in polluted ambient air (—■ 1—8 ppb) (see Chapter 11) taken with the Henry's law constant for HONO (Table 8.1) yield aqueous concentrations too low to contribute substantially to the aqueous-phase S(IV) oxidation.

For example, with a Henry's law constant for HONO of 49 M atm-1, a gas-phase concentration of 1 ppb would result in a solution-phase concentration of only 4.9 X 1CT8 mol L-1, compared to an anticipated H202 solution-phase concentration of 10~4 mol L_1 at the same gas-phase concentration. The rate constants also favor the H202 reaction; at a pH of 3.0, that for oxidation of H202 is approximately a factor of 104 larger than that for reaction with HONO. Thus, the combination of concentrations and rate constants makes HONO unlikely to be a significant S(IV) oxidant in solution unless other oxidants such as 03 or H202 are absent.

It is interesting, however, that the HONO-HSO^ reaction has been shown to form a nitrene (HON:), which Mendiara and co-workers (1992) suggest could contribute to free radical formation in clouds and fogs.

Whether dissolved N02 contributes significantly to the S(fV) oxidation in solution is uncertain. As seen from the Henry's law constant in Table 8.1, N02 is relatively insoluble; thus at a gas-phase concentration of 10 ppb, the equilibrium aqueous-phase concentration is only 10"10 mol L However, Schwartz and co-workers (Schwartz, 1984a) have inferred from literature data that the rate constants for the N02 reaction with HSO^ andS03~ may be sufficiently large, 3 X 105 and 1 X 107 L mol-1 s_l, respectively, that the N02-S(IV) reaction could be significant. Littlejohn et al. (1993) suggest that the reaction is that of N02 with sulfite, leading to the formation of SO^ and SO^ radical anions. The N02-S(IV) reaction may also be catalyzed in solution by the presence of carbon particles (Schryer et al., 1983). Further support for an aqueous-phase oxidation of S(IV) by NOz comes from cloud chamber studies where significant sulfate production was observed when N02 was present even at concentrations as low as 5 ppb N02 (Gertler et al., 1984).

The reaction of N03 scavenged from the gas phase with S(IV) is discussed in Chapter 7 and, as seen there, may also be important (e.g., Chameides and Davis, 1983; Rudich et al., 1996, 1998).

g. Free Radical Reactions in Clouds and Fogs

As might be expected based on our knowledge of the gas-phase chemistry, there are a variety of free radicals in atmospheric aqueous systems as well and these too can participate in S(IV) oxidation in clouds and fogs. The free radicals arise either from absorption from the gas phase or, alternatively, from in situ production, largely from photochemical processes. As seen from the data in Table 8.5, uptake of OH and H02 from the gas phase is fast, with mass accommodation coefficients >4 X 10~3 for OH and >0.2 for H02 on liquid water at 275 K (Hanson et al., 1992; Mozurkewich et al., 1987). Uptake of species such as H202, which can photolyze to form OH, is also fast.

In addition, there are many photochemical processes in clouds and fogs that can produce reactive species such as peroxyl radicals, singlet oxygen, 02('A ), OH, H02, and H202. Thus, a variety of studies have detected the formation of such species upon irradiation of rainwater, cloudwater, and fogwater (e.g., Faust and Allen, 1992, 1993; Zuo and Hoigné, 1993; Faust et al., 1993; Anastasio et al., 1994; Faust, 1994; Arakaki et al., 1995; Arakaki and Faust, 1998). The actual reactions leading to the formation of these oxidants are not well established. Suggested mechanisms include the reaction of organics to form superoxide ion, O^.

Zafiriou (1983), for example, suggested that absorption of light by organics, followed by intersystem crossing (ISC) to the triplet state (T) as described in Chapter 3, could occur. The subsequent reaction of the organic in a triplet state with 02 could then give 02 :

Organic chromophore (S0) +hv^> Chromophore (S,)

While energy transfer between triplets and 02 is well known, whether charge transfer can occur as shown in (2f) is not clear.

Superoxide ion and H02 are closely coupled via reaction (22):

0 0

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