Oxidations at Sulfur Centers by Aqueous Hypochlorous Acid and

Article pubs.acs.org/IC

Oxidations at Sulfur Centers by Aqueous Hypochlorous Acid and Hypochlorite: Cl+ Versus O Atom Transfer Ying Hu,†,‡ Guangyuan Xie,† and David M. Stanbury*,‡ †

College of Chemical Engineering, China University of Mining and Technology, Xuzhou 221116, People’s Republic of China Dept. of Chemistry and Biochemistry, Auburn University, Auburn, Alabama 36849, United States

‡

S Supporting Information *

ABSTRACT: Sulfur-containing compounds are known to be susceptible to oxidation by aqueous HOCl, but the factors affecting the rates of these reactions are not well-established. Here we report on the kinetics of oxidation of thiosulfate, thiourea, thioglycolate, (methylthio)acetate, tetrathionate, dithiodiglycolate, and dithiodipropionate at 25 °C and 0.4 M ionic strength. These reactions obey the general rate law −d[OCl−]/dt = (kOCl−[OCl−] + kHOCl[HOCl])[substrate] with some exceptions: tetrathionate and the two disulfides undergo rate-limiting hydrolysis at high pH, and dithiodiglycolate has an additional term in the rate law that is second order in [substrate]. The reactions of HOCl are believed to have a Cl+ transfer mechanism, and in the case of thiosulfate the rate of hydrolysis of the ClS2O3− intermediate was determined. In the case of thiourea evidence was obtained for thiourea monoxide as a long-lived product. It is shown that sulfite and species with terminal sulfur atoms have kHOCl values in the vicinity of 1 × 109 M−1 s−1, while SCN− and thioethers react somewhat more slowly; tetrathionate, trithionate, and disulfides react much more slowly. Comparison of the rate constants with those for oxidation of these sulfur substrates by H2O2 and [Pt(CN)4Cl2]2− shows that HOCl reacts a few orders of magnitude more rapidly than [Pt(CN)4Cl2]2− and ∼9 orders of magnitude more rapidly than H2O2. Many of the kHOCl values are leveled by the high electrophilicity of HOCl. It is proposed that the kOCl− values correspond to oxygen-atom transfer mechanisms, as supported by LFERS (linear free energy relationships) relating these rate constants to those for reactions of H2O2 and [Pt(CN)4Cl2]2−.

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INTRODUCTION Hypochlorous acid is an excellent aqueous oxidant for many reasons.1 Its cost is low, its reduction product is often the innocuous species Cl−, and its reactions tend to be relatively fast. Thermodynamically it is a fairly strong oxidant, with E°(HOCl, H+/Cl−, H2O) = 1.48 V; that is, it is stronger than O2 (E°(O2, 4H+/2H2O) = 1.23 V) but weaker than H2O2: (E°(H2O2, 2H+/2H2O) = 1.76 V).2,3 Kinetically HOCl can be much more reactive than either O2 or H2O2; for example, methionine is oxidized 1 × 1011 times faster by HOCl than by H2O2,4,5 and it has no perceptible direct reaction with O2. HOCl also has important roles as an intermediate in reactions of other oxidants: it can serve as a chain carrier in reactions of ClO2·,6,7 it is an essential component of chlorite oscillators,8 and it influences the isotopic product distribution in the reaction of ClO3− with SO32−.9 Moreover, it is often the predominant reactive species in solutions of Cl2. Extensive studies from the Margerum group have shown that HOCl tends to react with nucleophiles through a Cl+ transfer mechanism:10−15 HOCl + nuc− → nucCl + OH− nucCl + H 2O → nucOH + HCl

chlorinated species have been detected. Further support comes from the excellent LFER that relates the second-order rate constants for Cl+ transfer to the Swain−Scott nucleophilicities (n).11 In alkaline media OCl− is also an effective oxidant although less reactive than HOCl; at this point there is no literature consensus as to whether OCl− reacts through Cl+ or O atom transfer.11 Oxidations of many sulfur-containing species by HOCl/ OCl− occur readily. Nevertheless, HOCl can be extremely selective in its reactivity with these species, with reported rate constants ranging from 1 × 102 M−1 s−1 for S4O62− to 4.8 × 109 M−1 s−1 for HS−.16,17 Rate constants have also been reported for reactions of HOCl with S3O62−,18 (SCH2CH2CO2H)2,19 SO32−,12 cysteine,4,19 methionine,4,20 and SCN−.21 Despite these advances we still lack a clear understanding of the factors that influence the rates. Further complicating the matter, in some cases there are mechanistic ambiguities, and in some others the data are limited to a single pH. Two aspects of HOCl chemistry need to be stated at the outset: the species is a weak acid with pKaHOCl = 7.4 at 1 M ionic strength,22,23 and it readily undergoes comproportionation:24

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HOCl ⇌ OCl− + H+ pK a HOCl = 7.4

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Direct evidence for this Cl+ transfer mechanism has been obtained for CN−, SO32−, NO2−, and amines, where the © 2017 American Chemical Society

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Received: December 30, 2016 Published: March 14, 2017 4047

DOI: 10.1021/acs.inorgchem.6b03182 Inorg. Chem. 2017, 56, 4047−4056

Article

Inorganic Chemistry HOCl + Cl− + H+ ⇌ Cl 2 + H 2O

K = 9.6 × 102 M−2

quartz cell was 10 mm. Some of the experiments on the reaction of S2O32− with HOCl in acetate buffers were monitored at 250 nm to be consistent with a prior study. Oxidations under alkaline conditions by HOCl/OCl− were followed at 292 nm (the absorbance maximum of OCl−) under flooding conditions with at least a sevenfold molar excess of reductants over HOCl/OCl−; because of its very rapid kinetics the reaction with TGA was investigated at very low concentrations with [OCl−]0 = [TGA]0. Unless indicated otherwise, all kinetic data were obtained at 0.4 M ionic strength with NaClO4 as the inert electrolyte, and all concentrations specified refer to the mixed solutions. Five replicate measurements of the kinetic traces were obtained under each set of conditions to check the reproducibility of reactions. The nonlinear kinetic curves at 292 nm were fitted with an exponential decay function except for the reaction of HOCl/OCl− with TGA, where second-order kinetic fitting was employed. Mean kobs values were used in further analysis. Fits of the dependence equations kobs = f([reductant], [OH−]) were performed with the Prism 5 software package, weighting the data as 1/(kobs)2. Experiments designed to reveal the behavior of reactive intermediates in the HOCl/OCl− + thiosulfate reaction were performed with a pH indicator (2,4-dinitrophenol) in unbuffered solutions and monitored at 410 nm. Both reactant solutions were prepared in water and adjusted to pH ≈ 4.2 with H2SO4 before mixing in the stopped-flow instrument. Note that the indicator was added into only the reductant solutions (thiosulfate) because of its slow decomposition in the presence of HOCl. A high-performance liquid chromatograph (HPLC; Thermo Scientific UltiMate 3000) combined with a Phenomenex Gemini C18 column was applied for identifying products of HOCl/OCl− oxidation. The following conditions were used in studying the HOCl/ thiosulfate reaction: V(acetonitrile)/V(5 mM tetrapropylammonium hydroxide, pH adjusted to ∼7 with H3PO4) = 16:84, flow rate 1.0 mL/ min, detection wavelength 230 nm. For study of the HOCl/TU reaction the conditions were: V(H2O pH 3)/V(methanol)/V(acetonitrile) = 68:6:26, flow rate 0.4 mL/min, detection wavelength 214 nm.

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HOCl is typically many orders of magnitude more reactive than OCl−, which causes the net reaction rates to be strongly dependent on pH. The comproportionation reaction can also be a confounding factor, but it can be avoided by working in chloride-free media that are not too acidic. Here we report studies on the reactions of S2O32−, thiourea (TU, SC(NH2)2), thioglycolate (TGA, HSCH 2CO 2−), (methylthio)acetate (MTA, CH3SCH2CO2−), tetrathionate (S4O62−), dithiodiglycolate (DTDGA, (SCH2CO2−)2), and dithiodipropionate (DTDPA, (SCH2CH2CO2−)2). The S2O32− reaction reveals information about OS2O32−, a species believed to be an intermediate in many reactions of S2O32−. Likewise, the TU reaction generates an intermediate believed to be TUO (thiourea, S-oxide), which is also a proposed intermediate in many TU reactions. The TGA reaction was selected as a representative example of a thiolate to avoid the potential chlorination at the amine site present in the prior thiolate example, cysteine. Similarly, our study of MTA was undertaken to obtain data on a thioether that lacks the potential amine oxidation that could occur with methionine. Tetrathionate was selected for study because of its unusually low reported reactivity with HOCl. DTDGA and DTDPA were investigated as examples of disulfides, since the prior work on DTDPA was conducted at only a single pH value. With these new results we are now able to identify clear trends in the reactivity of sulfur compounds with HOCl.

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EXPERIMENTAL SECTION

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Materials. American Chemical Society certified grade or better commercially available reagents were used in these experiments, including NaOH (Sigma-Aldrich), acetic acid (Sigma-Aldrich), anhydrous sodium acetate (Sigma), sodium carbonate (Fisher Chemicals), sodium bicarbonate (J. T. Baker Chemical), sodium thiosulfate pentahydrate (Fisher Chemicals), potassium tetrathionate (Aldrich), 2,4-dinitrophenol (DNP, Aldrich), thiourea (Mallinckrodt Chemical Works), thiourea dioxide (Aldrich), formamidine disulfide (Tokyo Chemical Industry), (methylthio)acetic acid (Alfa Aesar), sodium thioglycolate (98%, ACROS Organics), dithiodiglycolic acid (Sigma), 3,3′-dithiodipropionic acid (Sigma-Aldrich), and yellow mercuric oxide (Mallinckrodt Chemical Works). Thiourea was purified by recrystallization at least twice from hot water. Dithiodiglycolic acid was recrystallized with hot acetone followed by air drying. NaOH stock solutions (∼1 M) were standardized by titration with potassium hydrogen phthalate. Chloride-free hypochlorous acid was prepared by a modification of a literature procedure:25 chlorine gas prepared from the reaction of KMnO4 with HCl was transferred into an aqueous suspension of HgO, followed by distillation of the mixture at reduced pressure. After alkalinization the solution was analyzed by UV/vis spectroscopy (λmax = 292 nm with ε = 350 M−1 cm−1).10,26,27 These solutions can be stored for one month at low temperature. All solutions were prepared with purified water with a specific resistivity of 18.2 MΩ cm at 25 °C from an Ultrapure water purification system. Working solutions of the reducing agents, such as thiosulfate, dithiodiglycolic acid, and 3,3′-dithiodipropionic acid, were prepared freshly in water to prevent decomposition at unfavorable pH values before usage. Instruments and Methods. An Agilent HP 8453 diode array spectrophotometer with quartz cells having calibrated 1 cm path lengths was used for collecting spectra of all compounds and determining the concentration of HOCl. Kinetic measurements for oxidation by HOCl/OCl− were performed at 25 ± 0.1 °C on a Hi-Tech SF-51 stopped-flow spectrophotometer equipped with a thermostat apparatus and using Olis 4300 data acquisition and analysis software. The light path of the

RESULTS AND DISCUSSION Reactions of simple substrates with hypochlorite in alkaline solutions typically obey a two-term rate law of the type rate = (kHOCl[HOCl] + k OCl−[OCl−])[substrate]

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When the pH is well-above the pKa of HOCl this rate law can be expressed as −d[OCl−]/dt ⎛ k ⎞ HOClK w − −⎟[substrate][OCl ] k = ⎜⎜ HOCl + OCl ⎟ [OH−] ⎝ Ka ⎠

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and when the concentration of substrate is in large excess over hypochlorite this leads to pseudo-first-order kinetics with ⎛ k ⎞ HOClK w −⎟[substrate] kobs = ⎜⎜ HOCl k + OCl ⎟ [OH−] ⎝ Ka ⎠ or

⎛ k ⎞ HOClK w −⎟ kobs/[substrate] = ⎜⎜ HOCl k + OCl ⎟ [OH−] ⎝ Ka ⎠

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It is usually the case that kHOCl is several orders of magnitude greater than kOCl−, so that relatively slow rates are approached at high pH. This general description applies to the substrates discussed below, except that modifications are made when the substrate is significantly basic and extremely reactive (TGA) or undergoes slow base hydrolysis (S4O62−, DTDGA, and DTDPA). 4048


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Inorganic Chemistry Reaction with Thiosulfate. In a recent report on the reaction of ClO2· with S2O32− it was argued on the basis of the Swain−Scott relationship that the reaction of HOCl with S2O32− should have a rate constant in the vicinity of 1 × 109 M−1 s−1.7 A similar value can be derived from the results of fitting a 10-step mechanism to the reaction of ClO2− with S2O32−.28 However, an earlier publication by Varga et al. on the HOCl/S2O32− reaction claimed a much smaller rate constant of 3.9 × 103 M−1 s−1.29 Two potential sources of concern with the work of Varga et al. are that the experiments were performed (1) with acetate buffers near pH 4 and (2) with a mixing device having a relatively long dead time (∼10 ms). The concern with the acetate buffers arises because such buffers are known to react with HOCl to form the reactive species acetyl hypochlorite.30 The concern about the dead time is that if the reaction was performed at pH 4 and it had a rate constant of 1 × 109 M−1 s−1 then the reaction would be complete well before the first data points in the kinetic traces. Here we report experiments designed to probe these concerns, aided by use of a stopped-flow instrument having a superior dead time (∼2 ms) and with attention paid to buffer effects. As shown in Figure 1A, when 2 mM S2O32− is mixed with 4 mM HOCl at pH 3.9 (0.2 M acetate buffer) the absorbency at

we examined the same reaction in a KH2PO4/K2HPO4 buffer at pH 6.5, and no obvious variation of absorbance at 250 nm could be detected. This latter observation can be understood if the HOCl/S2O32− reaction at pH 6.5 occurs within the dead time of our apparatus. Additional support for this interpretation is provided below, and it implies that the initial reaction of HOCl with S2O32− at pH 3.9 also occurs within the dead time; the absorbance changes observed at pH 3.9 thus arise from further slower reactions of the initial products. Although we do not propose a mechanism here to model the absorbance changes at pH 3.9, we do note that the absorbance changes at 292 nm should not be attributed to OCl−, since this species has negligible absorbance at pH 3.9. We found that it was possible to observe the reaction of hypochlorite with thiosulfate by performing stopped-flow experiments under highly alkaline conditions with excess thiosulfate and monitoring the absorbance at 292 nm, the absorbance maximum of OCl−. Under these conditions the absorbance decays with exponential kinetics and with an absorbance change corresponding to complete consumption of the OCl−, indicating a simple pseudo-first-order loss of OCl−. A detailed kinetics study of the HOCl/OCl−−thiosulfate reaction was performed by employing different concentrations of thiosulfate and hydroxide ion while maintaining at least an eightfold excess concentration of thiosulfate over OCl−. The dependence of kobs on the thiosulfate concentration was studied with 0.35 mM OCl− and 0.398 M hydroxide ion by varying the thiosulfate concentration from 3 to 25 mM. These results are displayed as a plot of kobs versus [S2O32−] in Figure 2

Figure 2. Thiosulfate concentration dependence of kobs for the reaction of thiosulfate with HOCl/OCl− at high pH. [S2O32−] = 3−25 mM, [OCl−]0 = 0.35 mM, [OH−] = 0.398 M, μ = 0.4 M (NaOH + NaClO4), T = 25 °C. Figure 1. Measured kinetic curves at 250 and 292 nm for the HOCl/ thiosulfate reaction in acetate buffers. [S2O32−]0 = 2 mM, [HOCl]0 = 4 mM. (A) pH = 3.91, [NaOAc] = 0.2 M, [NaClO4] = 0.3 M; (B) pH = 3.76, [NaOAc] = 0.02 M, [NaClO4] = 0.48 M.

(data in Table S1), the linearity of which implies that the reaction is also first-order with respect to thiosulfate. The hydroxide ion dependence was investigated in the range from 0.071 to 0.398 M with 0.35 mM OCl− and 5 mM thiosulfate. Figure 3 (data in Table S2) is a plot of kobs/[S2O32−] versus 1/ [OH−] that demonstrates an excellent linear relationship with a positive intercept. The data in Figures 2 and 3 are consistent with eq 8. Thus, values of kHOCl = (2.31 ± 0.05) × 109 M−1 s−1 and kOCl− = (2.26 ± 0.12) × 103 M−1 s−1 are calculated from the slope and intercept of the linear fit in Figure 3 by use of pKaHOCl = 7.4 for HOCl.22,23 Our value for kHOCl is in good agreement with our prior estimate based on the Swain−Scott relationship,7 and it is also in good agreement with value that can be derived from the results of fitting a 10-step mechanism to the reaction of ClO2− with S2O32−.28 The extremely high rate constant for thiosulfate

250 nm (λmax for HOCl) is characterized by an initial apparent increase in the first 0.1 s and followed by a decreasing trend with a small “bump” appearing at ∼1 s; a “bump” is also observed at ∼1 s in data obtained at 292 nm. These results at 250 nm disagree with those of Varga et al. obtained under the same conditions, where an increasing absorbance was observed between 1 and 5 s. Figure 1B shows analogous data obtained at a 10-fold lower buffer concentration (0.02 M), where the 1 s “bump” is no longer detected. From these results we infer that the data obtained by Varga et al. were affected by both the mixing apparatus and the acetate buffer. In a further experiment 4049


Article

Inorganic Chemistry ClS2 O3− + S2 O32 − → S4 O6 2 − + Cl−

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We investigated these aspects of the thiosulfate/HOCl reaction by using an established pH indicator method.13 Equimolar solutions (0.257 mM) of HOCl and thiosulfate (both adjusted to pH 4.060 ± 0.010) were mixed in the stopped-flow apparatus; the thiosulfate solution also contained 0.0118 mM 2,4-dinitrophenol (DNPH, pKa = 4.11) as an indicator. The detection wavelength was set at 410 nm to record the concentration change of the conjugate base, DNP−. Figure 4 shows the absorbance changes upon mixing. The Figure 3. Plot of kobs/[S2O32−] vs 1/[OH−] for the reaction of thiosulfate with HOCl/OCl−. [OH−] = 0.071−0.398 M, [OCl−]0 = 0.35 mM, μ = 0.4 M (NaOH + NaClO4). (●) Second-order rate constants that are calculated by dividing the observed pseudo-firstorder rate constants by [S2O32−]. () the linear least-squares fit of the data with eq 8, kHOCl = (2.3 ± 0.05) × 109 M−1 s−1, kOCl− = (2.26 ± 0.12) × 103 M−1 s−1.

oxidation by HOCl implies that the reaction of thiosulfate and HOCl at pH 4−6 should be complete within the dead time of the stopped-flow instrument, consistent with our observations under these conditions. We thus infer that the 1 × 106-fold lower value for kHOCl reported previously for reactions in acetate buffers was derived from reaction features occurring after actual reaction of HOCl with S2O32−.29 That kHOCl is 6 orders of magnitude greater than kOCl− is typical of hypochlorite chemistry, as is seen, for example, in the analogous reaction of sulfite with OCl−/HOCl.12 High-performance liquid chromatography was utilized for the rapid determination of the products of the HOCl/OCl− + S2O32− reaction. Experiments performed with 2.98 mM S2O32− and 0.56 mM OCl− at pH 12.7 showed that S4O62− is a major product, but the S4O62− concentration decreases to the detection limit as the product solution is aged (∼10 min). Yields more consistent with quantitative formation of S4O62− were obtained by adjusting the product solution pH to ∼9 immediately after mixing. These results are as expected, given that S4O62− is well-known to undergo base-catalyzed hydrolysis on this time scale.31 Neglecting this relatively slow product decomposition, the reaction of HOCl with excess S2O32− is thus given as HOCl + 2S2 O32 − → S4 O6 2 − + Cl− + OH−

Figure 4. Indicator (DNPH) observation of the HOCl−S2O32− reaction. Initial pH = 4.160 ± 0.003, [DNP]tot = 0.0118 mM, T = 25 °C. (A) [HOCl]0 = [S2O32−]0 = 0.257 mM; (B) 2[HOCl]0 = [S2O32−]0 = 0.514 mM; (C) [HOCl]0 = [S2O32−]0 = 0 mM. Red line: the fit of the data in A with eq 14, kh = 75 ± 0.5 s−1.

initial absorbance is higher than what is observed in the absence of S2O32− and HOCl, because reaction 10 occurs in the dead time at this pH, leading to an initial jump in pH; after the mixing, the absorbance decreases exponentially (line A in Figure 4), consistent with the release of H+ during hydrolysis of ClS2O3− (eq 11). The final absorbance is considerably below the initial value, consistent with the sum of reactions 10 and 11. Trace B in Figure 4 shows that there is no observable absorbance change when the reaction is performed with a twofold excess of thiosulfate; this can be understood if S2O32− reacts with OS2O32− as quickly as it is produced (eq 12) or if it reacts with ClS2O3− within the instrument dead time (eq 13). An experiment at pH 11.8 with equimolar OCl− and S2O32− and Tropaeolin O as the indicator (as suggested by Nagy et al.)32 showed quite small absorbance changes at 550 nm, which suggests that the hydrolysis of ClS2O3− is base-catalyzed. On the basis of the assumption that the hydrolysis of ClS2O3− at pH 4.1 obeys first-order kinetics, the integrated rate law can be derived, which leads to eq 14 as an expression of the time dependence of the absorbance. Here kh is the hydrolysis rate constant (eq 11; see Supporting Information). 1 A= (1/A 0 − 1/A∞)e−kht + 1/A∞ (14)

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As mentioned above, HOCl typically reacts with nucleophiles through a Cl+ transfer mechanism, followed by hydrolysis. In the case of thiosulfate this would correspond to reactions 10 and 11, with kHOCl being the rate constant for Cl+ transfer (reaction 10). HOCl + S2O32 − → OH− + ClS2O3−

kHOCl

ClS2 O3− + H 2O → Cl− + OS2 O32 − + 2H+

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kh

Fits of the data to eq 14 yield kh = 75 ± 0.5 s−1. This is lower than the value reported for hydrolysis of ClSO3− (270 s−1)13 but greater than the value for CH3SO2Cl (8 × 10−5 s−1).33 In an earlier report on the reaction of S2O32− with ClO2− a 10-step mechanism was developed, and one of the steps was the reaction of ClS2O3− with S2O32− (eq 13 above).28 The parameters in that kinetic model only lead to a lower limit of 60 M−1 s−1 for the rate constant of reaction 13 and thus are not inconsistent with our value of kh. A rate constant for reaction 13 considerably greater than 60 M−1 s−1 would also be sufficient

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If the second step is slower than the first, this reaction sequence will lead to a characteristic rise and then fall in pH when unbuffered solutions are used, as has been demonstrated in the reaction of HOCl with SO32−.13 Under conditions of excess thiosulfate, the net reaction 9 would be attained by the rapid reaction of OS2O32− or ClS2O3− with S2O32−: OS2 O32 − + S2 O32 − + 2H+ → S4 O6 2 − + H 2O

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Inorganic Chemistry for excess S2O32− to consume the ClS2O3− via reaction 13 wihin the instrument dead time as mentioned above. Reaction with Thiourea (NH2)2CS. It has been proposed that the hypochlorite−thiourea reaction should be a component of the mechanisms of the thiourea oxidation by both ClO2− and ClO2.34,35 However, to our knowledge there have been no reports on the kinetics of the direct reaction of hypochlorite with thiourea. Here we report stopped-flow kinetics measurements at high pH and product information based on both UV−vis spectra and HPLC analysis. Figure 5 shows kinetic traces at 292 nm that were observed after mixing 0.11 or 0.24 mM OCl− with 1 mM thiourea in 0.01

has a very low absorptivity, TUO2 has approximately half of the molar absorptivity of OCl− at 292 nm (Figure S3), while TUO3 has essentially no absorptivity at this wavelength.36 Formation of those species would lead to an absorbance decrease rather than the rise shown in Figure 5. The HPLC chromatogram of the product mixture (Figure S4) clearly reveals the formation of thiourea monoxide (TUO, (NH2)2CSO) by comparing the retention time with that reported in literature;37 moreover, the UV spectrum of this species obtained by HPLC with diodearray detection corresponds to the spectrum of the product mixture (Figure S5). Figure S5a also shows that TUO has an absorbance maximum at 260 nm and a significant but smaller absorbance at 292 nm. A Job plot analysis at 260 nm exhibits a maximum when the ratio [OCl−]/([OCl−] + [TU]) is 0.5, indicating a thiourea/OCl− stoichiometric ratio of 1:1 (Figure S6), consistent with formation of TUO. This Job plot shows a nonlinear decline of absorbance when the initial ratio [OCl−]/ [TU] exceeds 1; this nonlinearity is probably due to the oxidation of TUO to TUO2 or TUO3 in the presence of excess OCl−. Overall, these results imply that the net reaction under conditions of excess TU at high pH is described by eq 15: TU + OCl− → TUO + Cl−

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It is assumed that the kHOCl pathway leads to eq 15 through the usual sequence of Cl+ transfer and hydrolysis. As is discussed below, the kOCl− pathway is suggested to proceed through direct O atom transfer. Reaction with Thioglycolate. On the one hand, our preliminary data showed that the reaction of thioglycolate (TGA, HSCH2COO−), with HOCl/OCl− is too fast to be monitored under pseudo-first-order conditions with stoppedflow technology, even in alkaline solutions. On the other hand, by working with low equimolar concentrations of both reagents at high pH it was possible to observe reproducible absorbance decays. On the basis of the assumption that the reaction is firstorder with respect to both reagents, the rate law can be expressed by the following second-order kinetic equation when the initial concentrations of both reagents are equal:

Figure 5. Typical kinetic traces for the HOCl/OCl− + thiourea reaction. (black) Experimental data. (red) Exponential least-squares fit. [TU] = 1 mM, [OH−] = 0.01 M, kobs = 81 s−1 (for [OCl−]0 = 0.11 mM), kobs = 79 s−1 (for [OCl−]0 = 0.24 mM).

M sodium hydroxide solution. In both traces the initial absorbance is in agreement with the absorbance of the unreacted OCl−, indicating that the subsequent absorbance changes correspond to the rate of consumption of OCl−. Surprisingly, these absorbance changes are exponential rises, which implies that a product is generated that absorbs more than the reactants. Both the excellent exponential fitting and the nearly identical kobs values (79, 81 s−1) with different initial OCl− concentrations strongly imply that the reaction is firstorder with respect to OCl−. The kinetics of the reaction were examined as a function of thiourea and hydroxide concentration by keeping a constant [OCl−]0 = 0.07 mM under pseudo-first-order conditions. Increasing the thiourea concentration from 0.42 to 2.2 mM in 0.025 M sodium hydroxide leads to a linear increase of kobs, which indicates the reaction exhibits a first-order dependence on [TU] (Figure S1, Table S3). The hydroxide dependence was measured at six different concentrations (0.009−0.06 M) in 0.6 mM thiourea (Table S4). The plot of kobs/[TU] versus 1/ [OH−] (Figure S2) demonstrates an excellent linear relationship with a positive intercept, indicating that both HOCl and OCl− react with thiourea according to rate law 8. The values of the rate constants kHOCl = (3.28 ± 0.03) × 109 M−1 s−1 and kOCl− = (5.6 ± 0.3) × 103 M−1 s−1 are derived from the slope and intercept. This value for kHOCl is ∼1 × 106-fold greater than the values utilized for this process in simulating the reactions of ClO2− and ClO2 with TU; this raises the possibility that the models of these chain reactions require significant revision.34,35 Products with excess thiourea under kinetic conditions were investigated by various means. Formamidine disulfide (TU2, (NH2)(NH)CSSC(NH2)(NH)), thiourea dioxide (TUO2, (NH2)2CSO2), and thiourea trioxide (TUO3, (NH2)2CSO3) as major products can be excluded because of their relatively lower absorbances at 292 nm compared to that of OCl−: TU2

−d[OCl−]/dt = k 2obs[OCl−]2 = k 2obs[HOCl]tot [TGA] (16) −

Here, [HOCl]tot is the sum of [OCl ] and [HOCl]. A series of experiments with 0.18 mM concentrations of the two reagents from pH 9.8 to 13 shows that k2obs is highly sensitive to pH (Figure 6, Table S5). In this pH range the carboxylate

Figure 6. pH dependence of k2obs for the reaction of HOCl/OCl− with thioglycolate. [OCl−]0 = [TGA]0 = 1.8 × 10−4 M, μ = 0.4 M (NaOH/ NaClO4/NaHCO3/Na2CO3), T = 25 °C. (●) Pseudo-second-order rate constants; () the nonlinear least-squares fit of the data with eq 17, kHOCl = (6.9 ± 0.5) × 109 M−1 s−1, kOCl− = (9.6 ± 0.8) × 105 M−1 s−1. 4051


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Inorganic Chemistry group in TGA is fully ionized, and the thiol is partially ionized. Rate law 17 can be derived (see Supporting Information) under the assumption that only the thiolate form of TGA (−SCH2COO−) is reactive: ⎛ ⎞ K aHOCl [H+] ⎟ − k 2obs = ⎜⎜kHOCl HOCl + k OCl HOCl + + ⎟ Ka + [H ] Ka + [H ] ⎠ ⎝ K aRSH K aRSH

+ [H+]

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The solid line in Figure 6 shows the excellent fit achieved with eq 17 and pKaRSH = 9.96 at μ = 0.5 M,38 the fitted parameters being kHOCl = (6.9 ± 0.5) × 109 M−1 s−1 and kOCl− = (9.6 ± 0.8) × 105 M−1 s−1. Reaction with (Methylthio)acetate (CH3SCH2CO2−). This reaction is considerably slower than that of TGA, and hence it was investigated with a pseudo-first-order excess of MTA. The MTA dependence of kobs was investigated with a series of MTA concentrations from 1.5 to 12 mM in 0.1 M NaOH; it shows an excellent linear relationship between kobs and [MTA] (Figure S7 and Table S6), which implies a firstorder dependence of kobs on [MTA]. Similar experiments were performed over a [OH−] range from 0.025 to 0.1 M at 2 mM [MTA]. The results are displayed in Figure 7 as a plot of kobs/

Figure 8. Kinetic traces for OCl− + tetrathionate reaction at various S4O62− concentrations under alkaline solutions. [NaOH] = 0.01 M, [OCl−]0 = 0.34 mM.

and the fitted value of kS4O6 is 0.12 ± 0.01 M−1 s−1 (μ = 10−40 mM). A proposed mechanism is as follows: S4 O6 2 − + OH− → S2 O32 − + OS2 O32 − + H+

k hyd (19)

−

OCl + S2 O3

2−

−

→ Cl + OS2 O3

2−

(net) fast

(20)

OS2 O32 − + 3OCl− + H 2O → 2SO4 2 − + 3Cl− + 2H+

(net) fast

(21)

This mechanism leads to an overall stoichiometry of S4 O6 2 − + 7OCl− + 3H 2O → 4SO4 2 − + 7Cl− + 6H+ (22)

As a result, the relationship kS4O6 = 7khyd can be derived, leading to a value of 0.016 M−1 s−1 for khyd. This rate constant is within the range of tetrathionate hydrolysis rate constants reported previously.39 In an earlier study it was shown that attack of HOCl on tetrathionate with kHOCl = 100 M−1 s−1 is the rate-limiting step in neutral solutions.17 Evidently, at higher pH the attack of HOCl becomes slower, because HOCl is converted to the less-reactive form OCl−; simultaneously, base hydrolysis of tetrathionate becomes faster. Thus, the remarkably low reported value for kHOCl is consistent with our observation of rate-limiting hydrolysis at high pH. Also note that the reaction of OCl− with S4O62− is so slow at high pH that it does not prevent the accumulation of S4O62− in the reaction of OCl− with S2O32− (eq 9). Reactions with Dithiodiglycolate ( − O 2 CCH 2 SSCH 2 CO 2 − ) and 3,3′-Dithiodipropionate (−O2CCH2CH2SSCH2CH2CO2−) under Strongly Alkaline Conditions. DTDGA and DTDPA are soluble disulfides; DTDPA and other disulfides are known to decompose in alkaline solutions due to cleavage of the disulfide bonds, especially above pH 11;40 this behavior also occurs in disulfide analogues, such as tetrathionate and formamidine disulfide.31,41 Thus, experiments on the reactions of DTDGA and DTDPA with OCl− at high pH were performed by mixing neutral solutions of the disulfides with alkaline solutions of OCl−. In preliminary experiments on the reaction of OCl− with excess DTDGA, kinetic traces in 0.01 M OH− at 292 nm show that the reaction is close to zero-order (Figure S-9); at lower pH the kinetic traces display some curvature intermediate between zero- and first-order kinetics. This pH effect is attributed to a transition between rate-limiting base hydrolysis of DTDGA at high pH and direct oxidation of DTDGA by

Figure 7. Plot of kobs/[MTA] vs 1/[OH−] for the reaction of (methylthio)acetate with HOCl/OCl−. [MTA] = 2.0 × 10−3 M, [OCl−]0 = 2.0 × 10−4 M, μ = 0.4 M (NaOH + NaClO4). (●) Pseudosecond-order rate constants that are calculated by dividing kobs by [MTA]. () The linear least-squares fit of the data with eq 8, kHOCl = (5.99 ± 0.06) × 108 M−1 s−1, kOCl− = −11 ± 14 M−1 s−1.

[MTA] versus 1/[OH−]. The linear relationship indicates an inverse dependence of k obs /[MTA] on hydroxide ion concentration. A fit of the data in Figure 7 with eq 8 leads to a value for kHOCl of (5.99 ± 0.06) × 108 M−1 s−1 and a statistically insignificant value for kOCl−. Reaction with Tetrathionate. The reaction of OCl− with excess tetrathionate (S4O62−) at high pH occurs with a linear decay of absorbance at 292 nm, indicating pseudo-zero-order kinetics. Kinetic traces in 0.01 M NaOH are shown in Figure 8 for a range of tetrathionate concentrations. These traces are increasingly offset with increasing [S4O62−] because of the S4O62− absorbance, and they show that the rate increases as the tetrathionate concentration increases. Figure S8A shows that the rates depend linearly on [S4O62−] in 0.01 M [OH−], and Figure S8B shows that the rates depend linearly on [OH−] at 6 mM S4O62−. Thus, the overall rate law is −d[OCl−]/dt = k S4O6[S4 O6 2 −][OH−]

(18) 4052


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Inorganic Chemistry

These results indicate that the rate law under these conditions is

HOCl at lower pH. On the one hand, data obtained in 0.04 M OH− show that the zero-order rates are linearly dependent on [DTDGA] (Figure S10 and Table S8); on the other hand, experiments with [OH−] ranging from 0.04 to 0.4 M at 2.4 mM DTDGA show that the rates depend linearly on [OH−] but have a small intercept (Figure 9 and Table S9).

−

d[OCl−] = kHOCl[HOCl][DTDPA] dt k K w [DTDPA][OCl−] = HOCl OH− K aHOCl

(27)

Thus, HOCl oxidizes DTDPA with a rate constant kHOCl of (9.6 ± 0.2) × 105 M−1 s−1, while oxidation by OCl− is imperceptibly slow. A Cl+ transfer mechanism is inferred, as in eq 28: RSSR + HOCl → RS(Cl+) → SR + OH−

kHOCl

(28)

The RS(Cl )SR product should then transform to final products through hydrolysis and oxidation with HOCl. Pattison and Davies investigated the reaction of HOCl with DTDPA at pH 7.3 (0.1 M phosphate buffer) and 20 °C;19 they obtained an effective rate constant of (1.6 ± 0.6) × 105 M−1 s−1, which is in reasonable agreement with the value of 5.3 × 105 M−1 s−1 that can be derived at pH 7.3 from our value of kHOCl. An unusual result is obtained in the reaction of DTDGA with HOCl. As Figure 10 shows, the plot of kobs versus [DTDGA] at pH 9.89 shows obvious upward curvature (data in Table S12). +

Figure 9. Hydroxide ion dependence in the HOCl/OCl− + DTDGA reaction at high pH. [DTDGA] = 2.4 mM, [OCl−]0 = 0.2 mM, [OH−] = 0.04−0.4 M, μ = 0.4 M (NaOH + NaClO4), T = 25 °C. (−d[OCl−]/dt)/[DTDGA] = kH2O + kOH−[OH−].

These results show that the rate law at high pH is −d[OCl−]/dt = k OH−[DTDGA][OH−] + kH2O[DTDGA] (23) −2

The fitted rate constants are kOH = (7.7 ± 0.1) × 10 M−1 s−1 and kH2O = (3.8 ± 0.3) × 10−3 s−1. A mechanism is proposed for the kOH− pathway that is analogous to the tetrathionate reaction described above: −

RSSR + OH− → RS− + RSOH −

k OH−

(24)

fast

(25)

−

RS + HOCl/OCl → products −

RSOH + HOCl/OCl → products

fast

Figure 10. DTDGA concentration dependence of kobs in the HOCl/ OCl− + DTDGA reaction under weakly alkaline conditions. Solid line is the result of fitting with eq 29. [OCl−]0 = 0.2 mM, [DTDGA] = 1.5−15 mM, NaHCO3/Na2CO3 buffer, pH = 9.89, μ = 0.4 M (NaHCO3/Na2CO3).

(26)

The origin of the kH2O term in rate law 23 is unclear, but it may be due to a minor contribution from the direct reaction of DTDGA with HOCl/OCl−. Experiments on mixtures of DTDPA with OCl− at high pH (∼13.6) show no obvious loss of absorbance at 292 nm on the stopped-flow time scale. These results imply that DTDPA has an extremely small rate of base hydrolysis, consistent with a published report that kOH− for DTDPA is 1.9 × 10−5 M−1 s−1 at pH 13.8 (40 °C).40 Reactions with 3,3′-Dithiodipropionate and Dithiodiglycolate under Weakly Alkaline Conditions. Unlike at high pH, in carbonate buffers around pH 10 the reactions of DTDPA and DTDGA with HOCl/OCl− are characterized by pseudo-first-order decays of the absorbance at 292 nm. Evidently, direct reaction of the disulfides with HOCl/OCl− dominates over base hydrolysis in these less-alkaline solutions. A detailed study of the kinetics of the reaction of HOCl/ OCl− with DTDPA was performed at pH values ranging from 9.56 to 10.56 and disulfide concentrations ranging from 1 to 7 mM. Figure S11 displays the plot of kobs versus [DTDPA] at pH 9.95, the linearity of which clearly demonstrates a first-order dependence on [DTDPA] (data in Table S10). The hydroxide dependence at 2.0 mM DTDPA is illustrated in Figure S12 as a plot of kobs/[DTDPA] versus 1/[OH−] (data in Table S11); this plot is strictly linear with an undetectably small intercept.

However, a more typical inverse dependence of kobs on [OH−] is displayed in Figure S13 (data in Table S13). These results can be accommodated by rate law 29 with kHOCl = 2.2 × 105 M−1 s−1 and kcat = 3.3 × 107 M−2 s−1: −d[OCl−]/dt = kHOCl[HOCl][DTDGA] + kcat[HOCl][DTDGA]2

(29)

The value of kHOCl for DTDGA is somewhat smaller than for DTDPA, which is opposite of the trend for the base hydrolysis rate constants. This difference can be understood, since the disulfides are nucleophiles in the Cl+ transfer reactions but electrophiles in the hydrolysis reactions. We suggest that the second term in the rate law reflects catalysis by the acyl hypochlorite of DTDGA, analogous to the catalysis of hypochlorite reactions by acetic acid.30 Similar catalysis might have been encountered in the DTDPA reaction if experiments had been performed at higher [DTDPA]. 4053


Article

Inorganic Chemistry Table 1. Measured Rate Constants for Reaction of HOCl/OCl− with Sulfur Compounds kHOCl, M−1 s−1

reactant

(2.31 ± 0.05) (3.28 ± 0.03) (6.9 ± 0.5) × (5.99 ± 0.06) N/Da (9.6 ± 0.2) × 2.2 × 105

S2O32− thiourea thioglycolate (methylthio)acetate S4O62− 3,3′-dithiodipropionate dithiodiglycolate a

kOCl−, M−1 s−1

× 10 × 109 109 × 108

comment ClS2O3− observed

(2.26 ± 0.12) × 103 (5.6 ± 0.3) × 103 (9.6 ± 0.8) × 105 N/Da N/D N/D N/D

9

105

zero-order at high pH third-order term

N/D: Not detected.

Table 2. Rate Constants for Reaction of Sulfur Compounds with Two-Electron Oxidants reactant S2O32− thiourea TGA HS− cysteine SCN− MTA methionine S3O62− S4O62− DTDPA DTDGA SO32−

kHOCl, M−1 s−1 2.3 3.3 6.9 4.8 1.2 2.3 6.0 7.0 6.6 1.0 9.6 2.2 7.6

× × × × × × × × × × × × ×

9a

10 109 a 109 a 109 g 109 i 107 k 108 a 108 l 103 o 102 p 105 a 105 a 108 q

kOCl−, M−1 s−1 2.3 5.6 9.6 2.7 1.9

× × × × ×

3a

10 103 a 105 a 104 g 105 i

2.3 × 104 q

kH2O2, M−1 s−1

kPt(IV), M−1 s−1

b

0.025 0.07b 18e 0.48h 15.2j 5.2 × 10−4 b

1.7 × 102 c,d

6 × 10−3 m

4n

0.202r

2.3 × 105 c,s

2.2 × 109 f 6.1 × 107 f 0.31c

a

This work. bReference 46. cReference 51. dReference 52. eReference 47. fReference 53. gReference 16. hCalculated from the data in ref 48. References 4 and 45. jReference 49. kReference 21. lReferences 4 and 20. mReference 5. nReference 54. oReference 18. pReference 17. qReference 12. rReference 50. sReference 55. i

■

sulfur compounds in Table 2, n values are available for SO32−, SCN−, S2O32−, HS−, and thiourea. With the exception of SCN− all of these species have values of kHOCl that are close to 1 × 109 M−1 s−1, which seems to be an upper limit for HOCl reactions. Thus, these species react at rates that are insensitive to n. SCN−, which reacts significantly more slowly, has the lowest value of n in this series. Although the Swain−Scott relationship is not very diagnostic for these sulfur compounds, evidence for Cl+ transfer is provided by the reactions of S2O32− (this work) and SO32−, where hydrolysis of the chlorinated products has been observed. Comparisons can also be made with the rate constants (kH2O2) for oxidation of some of the species in Table 2 by H2O2. It is immediately evident that for a given substrate the kH2O2 value is typically 8 to 11 orders of magnitude smaller than the kHOCl value. This reinforces the concept mentioned above that HOCl reacts much more rapidly than H2O2 despite its lower E° value, and it illustrates the value of chloroperoxidase enzymes in transforming H2O2 into the much more effective oxidant HOCl. The data in Table 2 also show that all of the substrates with reported kH2O2 values have kHOCl values in the range of (0.7−7) × 109 M−1 s−1 with the exception of SCN− (kHOCl = 2.3 × 107 M−1 s−1); thus, the correlation between the kHOCl values and the Swain−Scott parameters is also seen in the kH2O2 values. A qualitative extension of this trend is provided by the disulfides, which react slowly with HOCl and quite slowly with H2O2.44 One potentially complicating issue is that the peroxide reactions may not all have the same mechanism: some may react by addition (as in the SO32− reaction), while others may react by displacement of OH−.

GENERAL DISCUSSION A summary of the rate constants measured in this work is presented in Table 1. These results conform to the wellestablished trend that values of kHOCl are several orders of magnitude greater than the corresponding values of kOCl−. Another general trend is that the kHOCl values are all above 1 × 109 M−1 s−1 for the species with terminal sulfur atoms (S2O32−, (NH2)2CS, and −SCH2CO2−); the thioether MTA reacts somewhat more slowly, and the disulfides are much slower. This trend is explored more extensively in Table 2, where the values of kHOCl in Table 1 are combined with literature values for other sulfur compounds. It can be seen that HS− and cysteine join the group of highly reactive terminal-S compounds, but SCN− reacts noticeably more slowly. The fact that cysteine and thioglycolate have similar rate constants supports a mechanism in which halogenation occurs initially at the thiolate rather than the amine position in cysteine. The deviation for SCN− is explored below in comparisons with H2O2 and PtIV reactions. Table 2 also shows that methionine reacts similarly to (methylthio)acetate, confirming that thioethers react somewhat more slowly than terminal-S compounds. The entries for trithionate and tetrathionate show that these species react significantly more slowly than thioethers and disulfides. The low rate constants for the disulfides DTDPA and DTDGA provide a rationalization of the preferential N-halogenation of cystine and the glutathione dimer (GSSG) by HOCl.42,43 Finally, sulfite is seen to be highly reactive, even though the sulfur atom is three-coordinate. An excellent quantitative correlation of kHOCl values for a wide variety of nucleophiles has been demonstrated in an LFER with the Swain−Scott nucleophilicity parameter n.11 For the 4054


Article

Inorganic Chemistry [PtIV(CN)4Cl2]2− is another well-characterized two-electron oxidant, and it typically reacts through a Cl+ transfer mechanism. Values of its rate constants (kPt(IV)) for oxidation of six sulfur substrates are given in Table 2. It can be seen that the values of kPt(IV) are generally intermediate between the corresponding kHOCl and kH2O2 values. There is a strong linear correlation between the log(kPt(IV)) and log(kH2O2) values (Figure S14, slope = 2.1), which indicates that the factors determining nucleophilicity are the same in both sets of reactions. This correlation confirms again that SCN− is a relatively weak nucleophile despite having a terminal S atom. Overall, the relative lack of discrimination shown in the HOCl reactions is simply a consequence of the high eletrophilicity of HOCl, which has the effect of leveling the rate constants for the more reactive substrates. It is proposed that the kOCl− terms correspond to mechanisms in which OCl− reacts with the substrates by a rate-limiting oxygen-atom transfer mechanism, rather than the Cl+ transfer mechanism that is characteristic of HOCl reactions: OCl− + nuc → nucO + Cl− k OCl−

Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS Acknowledgment is made to the Donors of the American Chemical Society Petroleum Research Fund for partial support of this research. This research was also supported by a grant from Fundamental Research Funds for the Central Universities (China; Grant No. 2015XKZD09) and Priority Academic Program Development of Jiangsu Higher Education Institutions (PAPD). We thank Prof. Q. Gao of China Univ. of Mining and Technology for helpful discussions and use of the HPLC and UV−vis instruments in his laboratories. We thank S. Lymar of Brookhaven National Laboratories for his helpful comments.

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(1) Deborde, M.; von Gunten, U. Reactions of chlorine with inorganic and organic compounds during water treatment - Kinetics and mechanisms: A critical review. Water Res. 2008, 42, 13−51. (2) Bard, A. J.; Parsons, R.; Jordan, J. Standard Potentials in Aqueous Solution; Marcel Dekker, Inc.: New York, 1985; pp 834. (3) Milazzo, G.; Caroli, S. Tables of Standard Electrode Potentials; John Wiley & Sons: New York, 1978; pp 421. (4) Armesto, X. L.; Canle L, M.; Fernandez, M. I.; Garcia, M. V.; Santaballa, J. A. First Steps in the Oxidation of Sulfur-Containing Amino Acids by Hypohalogenation: Very Fast Generation of Intermediate Sulfenyl Halides and Halosulfonium Cations. Tetrahedron 2000, 56, 1103−1109. (5) Richardson, D. E.; Regino, C. A. S.; Yao, H.; Johnson, J. V. Methionine Oxidation by Peroxymonocarbonate, a Reactive Oxygen Species formed from CO2/Bicarbonate and Hydrogen Peroxide. Free Radical Biol. Med. 2003, 35, 1538−1550. (6) Hu, Y.; Horváth, A. K.; Duan, S.; Csekö, G.; Makarov, S. V.; Gao, Q. Mechanism Involving Hydrogen Sulfite Ions, Chlorite Ions, and Hypochlorous Acid as Key Intermediates of the Autocatalytic Chlorine Dioxide-Thiourea Dioxide Reaction. Eur. J. Inorg. Chem. 2015, 2015, 5011−5020. (7) Pan, C.; Stanbury, D. M. Kinetics of the Initial Steps in the Aqueous Oxidation of Thiosulfate by Chlorine Dioxide. J. Phys. Chem. A 2014, 118, 6827−6831. (8) Rábai, G.; Orbán, M. General Model for the Chlorite Ion Based Chemical Oscillators. J. Phys. Chem. 1993, 97, 5935−5939. (9) Halperin, J.; Taube, H. The Transfer of Oxygen Atoms in Oxidation-Reduction Reactions. III. The Reaction of Halogenates with Sulfite in Aqueous Solution. J. Am. Chem. Soc. 1952, 74, 375−380. (10) Johnson, D. W.; Margerum, D. W. Non-Metal Redox Kinetics: A Reexamination of the Mechanism of the Reaction between Hypochlorite and Nitrite Ions. Inorg. Chem. 1991, 30, 4845−4850. (11) Gerritsen, C. M.; Margerum, D. W. Non-Metal Redox Kinetics: Hypochlorite and Hypochlorous Acid Reactions with Cyanide. Inorg. Chem. 1990, 29, 2757−2762. (12) Fogelman, K. D.; Walker, D. M.; Margerum, D. W. Non-Metal Redox Kinetics: Hypochlorite and Hypochlorous Acid Reactions with Sulfite. Inorg. Chem. 1989, 28, 986−993. (13) Yiin, B. S.; Margerum, D. W. Kinetics of Hydrolysis of the Chlorosulfate Ion. Inorg. Chem. 1988, 27, 1670−1672. (14) Margerum, D. W.; Gray, E. T., Jr.; Huffman, R. P. Chlorination and the Formation of N-Chloro Compounds in Water Treatment. In Organometals and Organometalloids; Brinckman, F. E., Bellama, J. M., Eds.; ACS: Washington, DC, 1978; pp 278−291. (15) Gray, E. T.; Margerum, D. W.; Huffman, R. P. Chloramine Equilibria and the Kinetics of Disproportionation in Aqueous Soution. In Organometals and Organometalloids. Occurence and Fate in the Environment; Brinckman, F. E., Bellama, J. M., Eds.; ACS: Washington, D.C., 1978; pp 264−277.

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As is shown in Figure 11, plots of log(kPt(IV)) and log(kH2O2) versus log(kOCl−) (data in Table 2) are linear with slopes of 2.7

Figure 11. LFERs of log(kPt(IV)) and log(kH2O2) vs log(kOCl−). Slope(Pt(IV)) = 2.7. Slope(H2O2) = 1.2.

and 1.2, repectively. The slope close to 1 for the H2O2 reactions supports the concept that both the OCl− and H2O2 sets of reactions have O atom transfer mechanisms, while the slope of 2.7 for the Pt(IV) reactions further supports the concept that Cl+ transfer reactions have greater nucleophilic discrimination than O transfer reactions.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.6b03182. Tabulated and plotted rate-law concentration dependence data, UV spectra, HPLC analysis data, TUO spectra, Job plots, tabulated rate-law pH dependence data, stopped-flow kinetic traces, derivation of equations 14 and 17 (PDF)

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REFERENCES

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

David M. Stanbury: 0000-0002-3892-9048 4055


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Oxidations at Sulfur Centers by Aqueous Hypochlorous Acid and

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