Oxidation of aqueous sulfur dioxide by peroxymonosulfate

Recent model calculations suggest that peroxymonosulfate may constitute a significant fraction of the total sulfur budget in remote tropospheric water droplets such as cloud, fog, and rain. However, little is known about the oxidation of dissolved SO_2 by peroxymonosulfate (HSO_5^-). We have found in aqueous solution that the rate of S(IV) oxidation is comparable to the rate of oxidation of S(IV) by hydrogen peroxide and that HSO_4^- is the only detectable oxidation product. We propose a mechanism in which the rate-determining step involves the acid-catalyzed decomposition of a peroxide-bisulfite intermediate to disulfate ion, S2_7O^(2-), and ultimately to sulfuric acid. The rate equation for this mechanism is -d[HSO_3^-]/dr = k_1(k_2/k_(-1))K_(a1){H^+}[HSO_5^-][S(IV+)]/(1 + (k_2/k_(-1){H^+})(k_(al) + {H+))), where k_1 = 1.21 x 10^6 M^(-1) s^(-1), k_2/k_(-1) = 5.9 M^(-1), k_1k_2/k^(-1) = 7.14 x 10^6 M^(-2) s^(-1), K_(a1) = 2.64 x 10^(-2) M at 5 ºC, and µ = 0.2 M. The activation parameters are ΔH_(k1) = 25.74 ± 0.77 kJ mol^(-1) and ΔS_(k1)^*= -88.1 ± 2.7 J mol^(-1) K^(-1)

Kinetics, mechanism and thermodynamics of the reversible reaction of methylglyoxal (CH_3COCHO) with sulfur (IV)

At pH ≤ 2 the following rate law for the formation of hydroxyacetylmethaneulfonate (HAMS) from methylglyoxal (MG) and S(IV) (H_2O•SO_2, SO_3^(2-))is obtained: d[HAMS]/dt = ((k_0α_1[H^+]/K_(a0)) + k_1α_1 + k_2α_2)[S(IV)][MG]_0)/(1 + K_d + K_d[H^+]/K_(a0)), where α_1 and α_2 are the fractional concentrations of HSO_3^- and SO_3^(2-), respectively; k_0, is the rate constant for the reaction of HSO_3^- with the carbocation aldehyde species (CH_3COC^+HOH); k_1, and k_2 are the rate constants for the reaction of unhydrated MG with HSO_3^- and SO_3^(2-), respectively; K_d is the dehydration constant of hydrated MG; and K_(a0) is the acid dissociation constant of the carbocation. At pH ≥4 the rate of formation of HAMS is determined by the rate of dehydration of the diol form of (hydrated) MG: d[HAMS]/dt = k_d[MG]/(l + K_d + _Kd [H^+]/K_(a0)), where k_d = k_w + k_H[H^+] + kOH[OH^-] + k_A[A] + kB_[B], and k_w is the intrinsic (water) rate constant; k_H and k_OH) are the specific acid and base rate constants; and k_A and k_B are the general acid (A) and base (B) rate constants. Between pH 2 and 4, biexponential kinetics are observed because, under our conditions, the rates of dehydration and of S(IV) addition become comparable. Over the pH range 0.7-7.0, the dissociation of HAMS follows the rate law: d[S(IV)]/dt = ((k_(-0)[H^+] + k_)-1) + k_(-2)K_(a3)/[H^+])K_(a4)[H^+][HAMS])/[H^+]]^2 + K_(a4)[H^+] + K_(a3)K_(a4)) where k_(-0) k_(-1), and k_(-2) are the reverse of the analogous forward rate constants defined above and K_(a3) and K_(a4) are the acid dissociation constants of the sulfortate anion and the sulfonic acid, respectively. Experiments to determine the effect of temperature on the rate (and equilibrium) constants indicate a marked effect of ΔS^* (and ΔS_(298)) on the relative magnitude of these constants

Kinetics and mechanism of the oxidation of aqueous hydrogen sulfide by peroxymonosulfate

The stoichiometry and mechanism of the oxidation of aqueous S(-II) by HSO_5^-, is similar to the oxidation of S(-II) by H_2O_2, but the rate of oxidation by HSO_5^-; is 3-4 orders of magnitude faster than the corresponding reaction with H_2O_2. A two-term rate law of the following form is found to be valid for the pH range of 2.0-6.3: -d[S(-II)]/dt = k_1[H_2S][HSO_5^-] + k_2K_(a1)[H_2S][HSO_5^-]/[H^+], where k_1 = 1.98 X 10^1 M^(-1) s^(-1),k2 = 1.22 X 10^4 M^(-1) s^(-1), and K_(al) = [H^+][HS^-]/[H_2S] = 2.84 X 10^(-8) M at 4.9 °C, µ = 0.2 M, and [S(-II)] = [H_2S] + [HS^-] + [S^(2-). At high pH and high [HSO_5^-]/[S(-II)] ratios SO_4^(2-) and H^+ formation are favored, whereas at low pH and low [HSO_5^-]/[S(-II)] ratios elemental sulfur (S_8) is favored as the principal reaction product. Peroxymonosulfate is a monosubstituted derivative of hydrogen peroxide that is thermodynamically more powerful as an oxidant than H_2O_2 and kinetically more reactive. These properties make HSO_5^- a potentially important oxidant in natural systems such as remote tropospheric clouds and also a viable alternative to H_2O_2 for the control of malodorous sulfur compounds and for the control of sulfide-induced corrosion in concrete sewers

Aldehyde-bisulfite adducts: prediction of some of their thermodynamic and kinetic properties

Stability constants (K_1) for the reaction of acetaldehyde and hydroxyacetaldehyde with NaHSO_3, determined spectrophotometrically in aqueous solution, were found to be (6.90 ± 0.54) x 10^5 M^(-1) and (2.0 ± 0.5) x 10^6 M^(-1) respectively, where K_1 (corrected for aldehyde hydration) = [RCH(OH)SO_3^-]/[RCHO][HSO_3^-] (µ = 1 0.2 M; 25 °C). Acid dissociation constants (pK_(a3)) of a series of α-hydroxyalkanesulfonate salts, RCH(OH)SO_3^-, were found to be 11.46 (CH_3-), 11.28 (H-)10.30(HOCH_2-),10.33 (C_(6-)H_5-),10.31 (CH_3CO-), and 7.21 (Cl_3C-)(µ = 0 M; 25 °C). Simple straight-line relationships were found to exist between Taft's σ^* parameter and a number of thermodynamic and kinetic properties of some aldehydes. K_1, K_(a3), and the rate constant for nucleophilic addition of SO_3^(2-) all increase linearly with σ^*. Carbonyl species such as halogenated derivatives of acetaldehyde, certain β- and γ-dicarbonyl aldehydes, and perhaps also some highly (halogen) substituted ketones, ie., all those species with ∑σ* ≥ ~1.5 (aldehydes) or ∑σ* ≥ ~2.5 (ketones), could be important S(IV) reservoirs

Kinetics and mechanism of the reduction of cobalt(II) 4,4',4",4"'-tetrasulfophthalocyanine by 2-mercaptoethanol under anoxic conditions

The kinetics and mechanism of reduction of cobalt(II) 4,4’,4’’,4'’’-tetrasulfophthalocyanine by 2-mercaptoethanol to yield cobalt(I) tetrasulfophthalocyanine and 2-hydroxyethyl disulfide under anoxic conditions were investigated and the following rate law was found: v = -d[Co^(II)TSP]_T/dt = k_2K_1[RSH]_T[Co^(II)TSP]_T{2(1 + ɑ_H+ /Ka_1,+ Ka_2/ɑ_H+)(l+ ɑ[RSH]_T)}, where k_2 is a rate constant for the rate-limiting electron-transfer step, and K_1 is the equilibrium constant for the complexation of a CoTSP dimer with thioethanol; K_(a1) and K_(a2) are the apparent acid dissociation constants of HOC_2H_4SH and HOC_2H_4S^-, respectively; ɑ is K_1/(1 + ɑ_(H+)/K_(a1)+ K_(a2)/ɑ_(H+)):ɑ_H+ is the hydrogen ion activity. A nonlinear least-squares fit of the experimental data to the above rate law gave k_2 = 228 ± 3.8 s^(-l) and K_1 = 117 ± 2.5 M^(-1) at 27 ºC at µ = 0.4 M