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Similar results were observed when Mn(I1) is added to the Fe(I1) system, the effect being most significant in the presence of some initially added Fe(II1). A series of typical results for the sulfite induced autoxidation of Fe(I1) as a function of Mn(I1) concentration are reported in Fig. 6. In this case the Mn(I1) effect seems to exhibit a square dependence on its concentration. Although the details of this synergistic effect are not well understood at present, the observed Mn(I1) dependence may suggest the participation of dimeric species. Again the reaction of Fe(II1) with Mn(I1) t o produce Fe(I1) and Mn(II1) can initiate the overall sulfite induced autoxidation of Fe(I1). ConclusionsThe work described in this report has resolved important aspects of the autoxidation of metal ions that is needed in order to complete the catalytic cycle for metal catalyzed autoxidation reactions of S(IV) oxides. The remarkable finding that sulfite can induce such oxidation processes and the resulting redox cycling of the metal ions may be of significance in accounting for the importance of such reactions in atmospheric oxidation processes. The relative concentration ratio of oxygen (or another oxidant such as H202 or 0,) and sulfite determines the oxidation state and the catalytic activity of the investigated metal ions. This concentration ratio will vary under atmospheric conditions depending on the oxygen and sulfite levels present, and as such have a dynamic effect on the overall autoxidation chemistry.The authors gratefully acknowledge financial support from the Deutsche Forschungsgemeinschaft, Bundesministerium fur Forschung und Technologie, and Commission of the European Communities. a) On leave from the Stopped-flow spectrophotometry was used to study the oxidation of S(IV) by H202 at 285 K and 298 K, respectively, in the pH range -0.3 to 13 in buffered aqueous solution under pseudo-first-order conditions (I = 0.5, 1.0 and 2.0 M NaC104, respectively). The reaction of HSOT with H202 is subject to generalacid catalysis whereas that of the s@,ion is not. The dependence of the experimental first-order rate constant on the concentration of the excess partner [H202] was studied in detail. -In the pH range -0.3 to 6 the rate is given by Eq. (1) (KRl and Ksl: acid dissociation constants of H302+ and SO2, respectively; [HI = proton activity).

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