S

10-2

10"6 10"5 10"4 10"3 10"2 10"1 1 10 Specific reaction volume (ml m"3)

FIGURE 8.12 Percentage of S02 converted to sulfate after a time interval At in small "haze" particles, fogs, and clouds as a function of the aqueous reaction volume; note that the time intervals for each one are different, reflecting how long they typically last in the atmosphere (adapted from Lamb et al., 1987).

areas due to their generation from erosion of the earth's crust. Iron has a particularly rich chemistry because it forms a variety of complexes with OH" and with various S(IV) aqueous forms (see, for example, Betterton, 1993; Brandt et al., 1994; and Millero et al., 1995). Figure 8.13 shows the calculated concentrations of various iron complexes as a function of pH in a solution containing f X 1(T6 M Fe3 + and 1 X 1(T5 M S(IV) at an ionic strength of 0.01. Hence elucidating

FIGURE 8.13 Calculated concentrations of iron species in aqueous solution for [Fe(III)] = 1 X 10"(> M, [S(IV)] = 1 X 10"5 M, and I = 0.01. The sulfur complexes are shown by the dotted lines (adapted from Martin, 1994).

the role of iron in the S(IV) oxidation has involved first understanding the nature of such complexes in solution. Further complicating the iron-catalyzed oxidation is that the mechanism changes from an ionic mechanism in the low-pH regime (0-3.6) to a free radical mechanism at higher pHs (4-7).

Table 8.6 shows two proposed mechanisms for the iron-catalyzed reaction at high acidities, in the pH range from 0 to 3.6. The recommended rate expression is given by (Martin, 1994):

Fe(III) refers to the sum of all three-valent iron in solution, i.e., Fe(III) = Fe3++ FeOH2 + + Fe(OH)2+ + Fe0HS03 + FeSO/+ etc. Thus the rate of the iron-catalyzed reaction in the low-pH region decreases with increasing [H + ], This means that it shows the behavior depicted in Fig. 8.9a; i.e., it is self-quenching. That is, as S(IV) is oxidized to the acid, the pH falls and the rate also decreases.

Despite the simplicity of the rate law implied by Eq. (O), the behavior of the kinetics is very sensitive to a variety of factors. Thus the reaction is inhibited not only by [H + ] but also by the ionic strength (/) of the solution, by both S(IV) and S(VI), and, at high pH, by organics. Martin (1994) gives expressions for the dependence of the rate constant k in Eq. (O) on I, S(IV), and S(VI). The effect of ionic strength may be due to effects on the stability of complexes, whereas the sulfate is thought to complex one or more of the catalytic species in the reaction. Because of these complexities, the rate expression in Eq. (O) only applies for [S(IV)] < f X 10~5, [Fe(III)] > 1 X 10"7, I < 0.01, and [S(VI)] < 1 X 10"4 mol L"1, where k = 6.0 s"1 (Martin, 1994).

TABLE 8.6 Some Proposed Mechanisms for the Catalyzed Oxidation of S(IV) in Aqueous Solutions

Hoffmann and Jacob (1984) Fe3+ + HSOj <-> FeSOj + H + FeSO, —> Internal redox

Addition of HSO3 Addition of 02 ——► Formation of products

Conklin and Hoffmann (1988) HSO3 <-> SOj- + H + Fe3 + + H20 «-> FeOH2 + + H + FeOH2+ + <-> H0Fe0S02 complex

H0Fe0S02 complex + 02 <-> 02 adduct acid-calaly/ed ,

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