Article pubs.acs.org/IC
Kinetics of the Benzaldehyde-Inhibited Oxidation of Sulfite by Chlorine Dioxide Changwei Pan,† Qingyu Gao,† and David M. Stanbury*,‡ †
College of Chemical Engineering, China University of Mining and Technology, Xuzhou 221116, People’s Republic of China Department of Chemistry and Biochemistry, Auburn University, Auburn, Alabama 36849 United States
‡
S Supporting Information *
ABSTRACT: There has been steady interest in the aqueous reaction of ClO2• with sulfur(IV) since the 1950s, and a wide variety of rate laws and mechanisms have been proposed. In neutral-to-alkaline media, the reaction is challenging to study because of its great rate. Here it is shown that benzaldehyde can be used as an additive to slow the reaction and make its rates more amenable to study. The rates can be quantitatively modeled by a mechanism that includes reversible binding of sulfur(IV) by benzaldehyde and a rate-limiting mixed second-order reaction of ClO2• with SO32−. The latter reaction occurs through parallel electron transfer from SO32− to ClO2• and oxygen-atom transfer from ClO2• to SO32−.
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INTRODUCTION The oxidation of sulfite (SO32−) to sulfate (SO42−) by ClO2• is known to be very fast in alkaline media, poses some interesting mechanistic questions, and is a component of the oscillating chlorite−sulfite reaction in acidic media.1 In an early report, Halperin and Taube investigated this reaction in strongly acidic media, where it is relatively slow; they found that the rate law is second-order in [ClO2•] and first-order in [SO2], and they used 18 O-tracer methods to show that direct oxygen-atom transfer occurs.2 Later, Suzuki and Gordon investigated the reaction between pH 8 and 13, using stopped-flow methods;3 in this pH range, the reaction is so fast that it approaches the limits of the stopped-flow instrument. Suzuki and Gordon reported that the form of the rate law is different from that in acidic media, being only first-order in [ClO2•] and [sulfur(IV)] with k = 8.6 × 105 M−1 s−1 at 10 °C and pH 10.0. They also performed stoichiometry studies, showing that ClO2− is the major chlorine-containing product. Subsequently, Huie and Neta obtained a rate constant at ∼25 °C of 2.7 × 106 M−1 s−1 at pH 11.4 by pulse radiolysis, which essentially confirmed the Suzuki and Gordon result.4 Merényi et al. measured a rate constant of 2.6 × 106 M−1 s−1 also by pulse radiolysis,5 and they used the dependence of the absorbance change on [SO32−] to measure the initial electron-transfer equilibrium constant: ClO2• + SO32 − ⇌ ClO2− + SO3•−
attained by Suzuki and Gordon. Horváth and Nagypál developed a 10-step mechanism that excludes reaction (1) and includes a major pathway in which atom transfer from ClO2• to SO32− occurs. These studies raise several points of concern. First, ClO2• is known to associate with ClO2− to form (ClO2)2•− with an association constant of 5.0 M−1.8 The experiments in which the electron-transfer equilibrium constant was measured (K1) were performed in 1 M ClO2− solutions; however, the formation of (ClO2)2•− was not taken into account. Despite this oversight, the reported value for K1 is rather close to that which is calculated (K1,calc = 3.0 × 103) from the component standard potentials [E°(ClO2•/ClO2−) = 0.936 ± 0.003 V; E°(SO3•−/ SO32−) = 0.73 ± 0.02 V].9 Given the close agreement between these two determinations of K1, it is puzzling that reaction (1) was excluded in favor of an atom-transfer process in Horváth and Nagypál’s mechanism. Conceivably, the difference arises from the differing pH ranges of the study. Second, the various mechanistic proposals have been heavily influenced by the stoichiometric data, and some of them are questionable. Specifically, at high pH, the reactions are so fast that the stoichiometric results might be affected by the relative rates of mixing and of reaction. Third, another potential source of error arises from the use of borate buffers in some of the experiments by Suzuki and Gordon: it is now known that boric acid forms a complex with sulfite.6 Fourth, there is a possibility that the studies in acetate buffer were affected by the reactivity of CH3COOCl; this species is known to be a powerful chlorinating agent and is produced from the reaction of acetate with HOCl, 10,11 with the latter being inferred as an intermediate in the ClO2•/SO32− reaction. In view of these concerns, we have developed a method to perform the ClO2•/SO32− reaction near neutral pH while
K1 = 2.1 × 103 (1)
ClO2−/sulfur(IV)
Frerichs et al. showed that the reaction can become oscillatory in a continuously stirred tank reactor, and they developed a mechanism that included the generation of ClO2• and its reaction with SO32−, as in eq 1;1 aspects of this work have been disputed by Huff Hartz et al.6 Most recently, Horváth and Nagypál investigated the ClO2•/sulfur(IV) reaction in acetate buffers (pH 3.6−4.5).7 They found that the reaction is much slower than that at higher pH, so they were able to obtain considerably greater kinetic detail than was © XXXX American Chemical Society
Received: November 29, 2015
A
DOI: 10.1021/acs.inorgchem.5b02770 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry maintaining rates that are slow enough to attain thorough mixing of the reactant solutions prior to chemical reaction. The method is based on the well-known rapid and reversible addition of sulfite to benzaldehyde.12 The results are described herein.
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EXPERIMENTAL SECTION
Reagents and Solutions. Stock solutions of ClO2• were prepared as described previously.13 Phosphate buffers were prepared from NaH2PO4 and Na2HPO4. Benzaldehyde (Aldrich, >99.5%) was used as supplied. All solutions were prepared with deionized water that had been purified by a Barnstead NanoPure system. Reactant solutions were sparged with argon and permitted to contact only glass, Teflon, and poly(ether ether ketone). ClO2• concentrations were determined spectrophotometrically at 360 nm (ε = 1200 M−1 cm−1).14 Instruments and Methods. Most of the kinetic studies were performed on a High-Tech SF-51 stopped-flow instrument in the 1cm-optical-path-length configuration. Absorbance/time data were recorded and analyzed with an OLIS 4300 data acquisition system. Reactions were performed at 25.0 °C and monitored at 360 nm. Values of kobs are the average of ∼8 shots. Experiments on the effect of variation of the benzaldehyde concentration were performed by using an Applied Photophysics model SX-20 MV stopped-flow spectrophotometer with a 150 W xenon lamp and a single-wavelength monochromator; the instrument was operated in the 1.0-cm-opticalpath-length configuration. pH measurements were performed at room temperature (22 °C) with a Corning 450 pH meter and a Mettler Toledo combination InLab micro pH electrode filled with 3 M NaCl. Capillary electrophoresis measurements were performed as described previously, with the alteration that the electrolyte consisted of 0.1 M boric acid, 5.0 mM sodium borate, 20 mM KNO3, and 0.01 mM hexadimethrine bromide.15 Concentrations of the analytes were determined by a comparison of the peak areas to standard calibration curves. The results are reported as the average of triplicate injections. Kinetic simulations were performed with the software Berkeley Madonna 9.0 using the Rosenbrock (stiff) method of numerical integration and with the parameter DTMAX set to 1.0 × 10−4.16 Fits of the rate law to the values of kobs were performed with the Prism 5.0 software package, weighting the data as (1/kobs2).17
Figure 1. pH dependence of kobs for the reaction of ClO2• with sulfur(IV) in the presence of 30 mM benzaldehyde. [ClO2•]0 = 0.2 mM, [sulfur(IV)]0 = 1.88 mM, and phosphate buffer, 25.0 °C. The solid line has a slope of 0.97 ± 0.06.
Figure 2. Dependence of kobs on [sulfur(IV)]. pH 7.53 (50 mM phosphate buffer), [ClO2•]0 = 0.19 mM, [benzaldehyde] = 30 mM, and 25.0 °C. The straight line has a slope of (4.48 ± 0.37) × 104 M−1 s−1 and an intercept of −9 ± 11 s−1.
of −9 ± 11 s−1. These results thus indicate a first-order dependence of kobs on [sulfur(IV)]0. The benzaldehyde dependence of the kinetics is shown as a plot of kobs versus [benzaldehyde] at pH 7.54 with [sulfur(IV)]0 = 1.88 mM in Figure 3 (data in Table S-3). The upper limit on
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RESULTS In a typical stopped-flow experiment, a 0.29 mM solution of ClO2• in 30 mM benzaldehyde was mixed with a 1.88 mM solution of sulfur(IV) in 30 mM benzaldehyde, with both solutions at pH 6.9 (50 mM phosphate buffer). Under these conditions, the bulk of sulfur(IV) is bound to benzaldehyde. Upon mixing of the ClO2• and sulfur(IV) solutions, the absorbance at 360 nm (where ClO2• has its absorbance maximum) decreases to zero with pseudo-first-order kinetics (kobs = 11.5 s−1). In contrast, when the same experiment is performed in the absence of benzaldehyde, the reaction is nearly complete within the instrument dead time (∼2 ms). The pH dependence of kobs was investigated between pH 6.9 and 7.86, maintaining the other conditions as in the above experiment with 30 mM benzaldehyde, and the results are displayed in Figure 1 as a plot of log kobs versus pH (data in Table S-1). The data obey a linear relationship with a slope of 0.97 ± 0.06, indicating an inverse dependence of kobs on [H+]: kobs = kobs′/[H+]. The dependence of kobs on the sulfur(IV) concentration was probed at pH 7.53 (50 mM phosphate buffer) with [ClO2•]0 = 0.19 mM, [benzaldehyde] = 30 mM, and [sulfur(IV)]0 varying from 1 to 4 mM. As is shown in Figure 2 (data in Table S-2), the data show a linear dependence of kobs on [sulfur(IV)] with a slope of (4.48 ± 0.37) × 104 M−1 s−1 and a negligible intercept
Figure 3. Dependence of kobs on [benzaldehyde]tot. pH 7.48 (50 mM phosphate buffer), [ClO2•]0 = 0.204 mM, [sulfur(IV)]tot = 1.88 mM, and 25.0 °C.
the benzaldehyde concentration is set by its solubility, which is only ∼50 mM. This figure shows that increasing concentrations of benzaldehyde lead to slower rates, but the exact relationship is not exactly a simple inverse dependence. Stoichiometry measurements were performed in order to obtain data under the conditions where benzaldehyde is present and clearly affecting the rates. Our capillary electrophoresis method was able to determine the concentrations of Cl−, SO42−, ClO3− and sulfur(IV) in the product solutions. ClO2− B
DOI: 10.1021/acs.inorgchem.5b02770 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry Scheme 1
Scheme 2
benzaldehyde as C6H5CH(OH)SO3−, over half of the free sulfur(IV) is in the form of SO32−, and the concentration of SO32− is inversely proportional to [H+]. These considerations imply that k in eq 4 is k1KaSIV/KBA, which leads to a value of 1.2 × 106 M−1 s−1 for k1. The close agreement between this result and the values of k1 reported by Suzuki and Gordon (8.6 × 105 M−1 s−1) and by Merényi et al. (2.6 × 106 M−1 s−1) provides strong support for Scheme 1 in eqs 5−8. Elaborations to Scheme 1 are required in order to explain the product distribution and the deviation from a simple inverse dependence on the benzaldehyde concentration. Qualitatively, the deviation of kobs from the inverse dependence on the benzaldehyde concentration can be attributed to the slow dissociation of sulfur(IV) from its benzaldehyde complex. The formation of ClO2− is assumed to occur as a result of electron transfer between SO32− and ClO2•, and the insignificant yield of S2O62− occurs because the SO3•− radical is oxidized by ClO2• faster than it reacts with itself. The formation of Cl− and ClO3− is a consequence of an oxygen-atom-transfer pathway in the reaction of SO32− with ClO2•. These ideas are developed in Scheme 2. The forward and reverse rate constants for association of benzaldehyde with sulfite (eq 7) are well-established; here we adopt the values reported by Olson et al.: kfBA = 0.71 M−1 s−1 and krBA = 1.5 × 10−4 s−1.12 Electron-transfer oxidation of SO3•− by ClO2• to form SO3 as in eq 11 has been proposed previously,3,24 and electron transfer from SO3•− is often suggested for reactions where SO32− undergoes one-electron oxidation: examples include the reactions of [Os(bpy)3]3+, [Ru(NH3)4phen]3+, [IrCl6]2−, [CuIII(H−3G4)]−, and [Fe(phen)3]3+.25−29 The reaction of ClO• with ClO2• (eq 12) is very fast, with a rate constant of kClOx = 7 × 109 M−1 s−1 having been reported by Wang and Margerum.30 Likewise, the reaction of HOCl with SO32− (eq 13) is very fast, with khypo = 7.6 × 108 M−1 s−1.31 Scheme 2 requires that Cl− and ClO3− be produced in equal yields, in agreement with our capillary electrophoresis results. This scheme provides no route to S2O62−, in agreement with our capillary electrophoresis results; S2O62− was identified as a minor product in a prior study of this reaction, albeit under quite different conditions: acetate buffer (pH ∼ 4.5), no benzaldehyde, and 5-fold higher concentrations of ClO2•.7 Scheme 2 could be modified by addition of the self-reaction of SO3•− radicals, which generates S2O62− in part;32 this would
was not detectable by this method because its retention time was the same as that of the system peak. When reaction mixtures were prepared with [ClO2•]0 = 0.23 mM and [sulfur(IV)]0 = 1.88 mM in 30 mM benzaldehyde at pH 7.36, the product mixture contained 0.065 ± 9% mM Cl−, 0.23 mM SO42−, and 0.041 ± 6% mM ClO3−. No peak attributable to S2O62− was detected. With the assumption that the chlorine balance consists of ClO2−, its yield was 0.124 mM. Note that ClO2− and sulfur(IV) can coexist for several minutes in these product mixtures despite the significant rate of reaction between ClO2− and SO32− at this pH; this arises because the bulk of sulfur(IV) is bound by benzaldehyde and thus removed from the reactive pool. Nevertheless, the reaction does occur, and so our observed yield of SO42− (0.23 mM) is somewhat inflated. On the other hand, the reaction of benzaldehyde with chlorous acid is only significant at low pH,18 and hence it does not interfere with our measurement of the ClO2− yield. These results are consistent with a description of the reaction as occurring through two parallel paths: 2ClO2• + SO32 − + H 2O → 2ClO2− + SO4 2 − + 2H+ (2) •
2ClO2 + 2SO3 −
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→ Cl +
2−
ClO3−
+ H 2O + 2SO4 2 − + 2H+
(3)
DISCUSSION To a first approximation, the kinetic data in Figures 1−3 conform to the rate law −d[ClO2•]/dt = k[sulfur(IV)][ClO2•] /([H+][benzaldehyde]) −5
(4)
−1
with k = (3.4 ± 1.1) × 10 M s . In this expression, [sulfur(IV)] refers to the total concentration of sulfur(IV), both free and bound to benzaldehyde. This rate law can be understood according to Scheme 1. The pKa of HSO3− is well-established and has a value of 6.85 under our conditions of 25 °C and μ = 0.1 M.19 The reversible reactions of sulfite with benzaldehyde (eqs 6 and 7) have been studied extensively;12,20−23 we adopt a value of 4.81 × 103 M−1 for KBA, as reported by Olson et al.12 The value for KaBAS is taken as 4.1 × 10−11 M from Betterton et al.21 Thus, within the pH range of these experiments (6.9−7.9), sulfur(IV) is largely bound to the C
DOI: 10.1021/acs.inorgchem.5b02770 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry
Contrary to Horváth and Nagpyál,7 we find no need to propose the second-order reaction of SO3ClO2•2− with itself and its reaction with ClO2•. We conclude that the equilibrium constant reported for reaction (1) by Merényi et al.5 was fortuitously in agreement with the value derived from E° values. Overall, the principal reaction features are the competing ratelimiting steps of electron and oxygen-atom transfer between the ClO2• and SO32− reactants.
require an adjustment of kSV (eq 11) but would not otherwise have any substantial effects. The oxygen-atom-transfer step in eq 10 (kOat) is strongly driven with ΔG° = −269 kJ mol−1, based on values of ΔfG° (kJ mol−1) = −744 (SO42−), − 486 (SO32−), 108 (ClO•), and 119 (ClO2•).33,34 We note that oxygen atom transfer from ClO2• has been proposed in several reactions of ClO2• including those with S2O32−,15 S5O62−,35 and styrene.36 It has also been proposed previously in the reaction of ClO2• with SO32−.3,7 Similarly, oxygen-atom transfer to SO32− from SO5•− occurs with a rate constant37 of 5.6 × 108 M−1 s−1 and ΔG° = −262 kJ mol−1 (calculated from ΔfG° values33,34,38,39). We assign a value of 4 × 105 M−1 s−1 to kOat in order the generate the observed yields of Cl− and ClO3−. This assignment leads to a ratio of 1:0.4 for the relative rate constants of electron and oxygen-atom transfer between ClO2• and SO32−. Scheme 2, supplemented with appropriate proton-transfer reactions of the phosphate buffer and reverse rate constants for all reactions as required by the principle of detailed balancing, leads to an adequate simulation of the experimental results. The full set of reactions and rate constants is described in the Supporting Information, along with a detailed description of the basis for the selected values. Simulations of this mechanism were performed by numerical integration of the associated set of differential equations. These simulations reproduce the observed and inferred yields of Cl−, ClO3−, and ClO2−, and they also reproduce the kinetics including the deviation from a strictly inverse dependence of kobs on the benzaldehyde concentration. This model might appear to conflict with the pulse-radiolysis report5 on the equilibrium constant for the electron-transfer reaction between ClO2• and SO32− (eq 1) because the parallel oxygen-atom-transfer path would prevent the electron-transfer step from reaching equilibrium. Even when the oxygen-atom pathway is disregarded, there are good reasons to believe that equilibrium was not attained. One is that the rate constant for reaction of ClO2• with SO32− is relatively slow (∼1 × 106 M−1 s−1), so the second-order decay of the SO3•− radicals (k = 4 × 108 M−1 s−1)32 will be competitive with their generation. Further, according to Scheme 2, the SO 3 •− radicals preferentially react with ClO2•, which is an even greater drain on the SO3•− concentration. Conceivably, the (ClO2)2•− radicals, which would be formed extensively from association of ClO2• with the 1 M ClO2− present in the pulse-radiolysis experiments, could react rapidly with SO32− and thereby ensure rapid electron-transfer equilibration; unfortunately, the original publication did not disclose the time at which the “equilibrium” absorbance measurements were made. If true equilibrium had been attained, then the derived equilibrium constant would likely agree with that derived from the relevant E° values within a factor of 5, given that the equilibrium constant for association of ClO2• with ClO2− is 5 M−1 and the molar absorptivity of (ClO2)2•− is probably quite similar to that of ClO2• at the observation wavelength (380 nm).8 These issues remain unresolved but do not contradict the mechanism in Scheme 2. In summary, the use of benzaldehyde as a sulfite “buffer” effectively slows the reaction of ClO2• with SO32− so that it is easily monitored with stopped-flow methods under mildly alkaline conditions. The results obtained support a simple mixed second-order rate law in agreement with some of the prior studies. A mechanism is inferred in which ClO2• reacts with SO32− primarily through electron transfer, but with a significant minor pathway being oxygen-atom transfer.
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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.5b02770. Tables of kobs values and the overall mechanism for kinetic simulation, including reactions, parameters, explanation, results of fitting, and references (PDF)
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS C.P. thanks the China Scholarship Council for a scholarship to perform portions of this research at Auburn University and the Fundamental Research Fund for the Central Universities (Grant 2015QNA19).
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REFERENCES
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DOI: 10.1021/acs.inorgchem.5b02770 Inorg. Chem. XXXX, XXX, XXX−XXX
