Kinetics and Mechanism of Sulfur(IV) Oxidation by Hydrogen Peroxide in Aqueous Phase: The Non‐Linear Parts of the pH‐Profile

Abstract

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).

Chaotic pH Oscillations in the Sulfur(IV)−Hydrogen Peroxide−Marble Flow System

“…The reaction time was informative for the first term of rate law 1R, and we could determine the numerical value of k 1 at different temperatures from our experimental data. k 1 ′ values for 10 and 4°C were calculated from the kinetics data available at 25 and 12°C 16 (Table 1).…”

“…The oxidation of SO 3 2- by H 2 O 2 is not competitive with reaction 1, and its contribution to the overall reaction can be neglected in the pH range of interest. The kinetics are described by eq 1R. ,− Equations 1 and 1R indicate that the reaction must be self-accelerating if no additional acid−base buffer is present: The reaction produces H + and its rate increases with increasing [H + ]. Shown in Figure are typical pH−time curves measured at different temperatures.…”

“…The kinetics are described by eq 1R. 4, [14][15][16] Equations 1 and 1R indicate that the reaction must be selfaccelerating if no additional acid-base buffer is present: The reaction produces H + and its rate increases with increasing [H + ]. Shown in Figure 1 are typical pH-time curves measured at different temperatures.…”

Uptake of gas‐phase SO2 in aqueous sulfuric acid: Oxidation by H2O2, O3, and HONO

“…There have been several studies of the reaction of S(IV) with H202 covering the pH range -0.3 to 12 [Hoffmann and Edwards, 1975;Penkett et al, 1979;Kunen et al, 1983;McArdle and Hoffmann, 1983;Lind et al, 1987;Jayne et al, 1990;Drexler et al, 1992;Lagrange et al, 1996]. To our knowledge, this study is the first to measure rates for oxidation of SO2 in concentrated sulfuric acid.…”

“…However, in this and a previous study of heterogeneous SO2(g ) q-H202(aq) reaction (using a different experimental technique [Jayne et al,1990]), we do not a have good explanation for the discrepancy in calculated and observed uptake rates of 802(g ). As was pointed out byDrexler et al [1992], H20 2 undergoes protonation which can affect the SO2-H202 reaction rate. Model calculations including protonation showed no evidence of such a reaction channel in our experimental data.…”