Non-Triiodide Based Autoinhibition by Iodide Ion in the Trithionate

May 20, 2010 - 6H2O f 3SO4. 2- + 8I- + 12H+ . It is also shown that the consumption of iodine is inhibited by iodide ion, but it cannot be simply expl...
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J. Phys. Chem. A 2010, 114, 6521–6526

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Non-Triiodide Based Autoinhibition by Iodide Ion in the Trithionate-Iodine Reaction Gyo¨rgy Cseko˝ and Attila K. Horva´th* Department of Inorganic Chemistry, UniVersity of Pe´cs, Ifju´sa´g u´tja 6., H-7624 Pe´cs, Hungary ReceiVed: March 24, 2010; ReVised Manuscript ReceiVed: May 3, 2010

The trithionate-iodine reaction has been studied spectrophotometrically in a slightly acidic medium at 25.0 ( 0.1 °C in acetate/acetic acid buffer monitoring the absorbance at 468 nm at the isosbestic point of the iodine-triiodide ion system in the absence and presence of initially added iodide ion at I ) 0.5 M ionic strength adjusted by sodium perchlorate. The stoichiometry of the reaction was found to be S3O62- + 4I2 + 6H2O f 3SO42- + 8I- + 12H+. It is also shown that the consumption of iodine is inhibited by iodide ion, but it cannot be simply explained via the fast equilibrium formation of triiodide ion. A five-step kinetic model with four fitted and fixed kinetic parameters is proposed and discussed in detail, on the basis of which all the most important characteristics of the measured kinetic curves are adequately explained. Introduction Although trithionate is thought to be at least as stable as tetrathionate in aqueous solutions, its oxidation by simple inorganic halogen-containing oxyanions are considerably less known. The reactions of the corresponding tetrathionate by iodine and chlorite, for example, have fascinating kinetics because they are autoinhibited by iodide ion1,2 and are autocatalytic with respect to hydrogen ion and hypochlorous acid as well as self-inhibited by chlorite ion, respectively.3 Moreover, the tetrathionate-chlorite reaction has recently become the subject of studying different types of spatiotemporal phenomena in connection with front propagation such as lateral instability4–6 and spatial bistability.7,8 Although the stoichiometry and kinetics of the reaction are rather complex,3,9,10 a simple stoichiometric equation along with a relatively simple rate equation was suggested to interpret the different spatiotemporal behaviors found experimentally.11 Our latest results, however, just showed that the formerly measured front velocity-concentration data pairs can be explained semiquantitatively with slightly modified stoichiometric and rate equations that also reflect to the complexity of the reaction.12 The complete quantitative description was not possible because it is well-known13 that tetrathionate is subject to decompose in alkaline solution into trithionate, sulfite, and thiosulfate. Unfortunately, no information has been available yet in the literature about the kinetics and mechanism of the trithionate-chlorite reaction. Moreover, systematic search among the articles published with the subject of trithionate has revealed the complete knowledge of only few oxidation reactions of trithionate. Since little is known about the redox reactions of trithionate but one may expect some analogies between that of the tetrathionate and trithionate due to their similar chemical structure, exploration of a less complex redox reaction of trithionate seems to be more preferable for a start to understand the chemistry of trithionate gradually. The reaction of iodine with trithionate may therefore be a good candidate since on the one hand the reaction can easily be followed by UV-vis spectroscopy and on the other hand the kinetics of the corresponding tetrathionate-iodine reaction is well-known.1,2 It was clearly demonstrated that the iodide inhibition mainly arose * To whom correspondence should be addressed: E-mail: horvatha@ gamma.ttk.pte.hu.

from the initiating fast equilibrium between tetrathionate and iodine despite the fact that trihalide ion is usually a less reactive species than the corresponding halogen molecule. The initiating equilibrium between the reactants resulted in the formation of S4O6I- and I- as a formal I+ transfer from iodine molecule to one of the inner sulfur atom of tetrathionate.1,2 The aim of our paper is, if it exists, to seek a general analogy between the redox reactions of tetrathionate and trithionate, as well as to unravel the kinetics and mechanism of the trithionate-iodine reaction, comparing it to that of the corresponding tetrathionate-iodine reaction. Experimental Section Materials and Buffers. The sodium trithionate was first prepared by the reaction of sodium thiosulfate with hydrogen peroxide as described previously.14 Unfortunately, the purity of the sodium trithionate was found to be less than 97%, therefore the relatively unpure solid substance was redissolved in a small amount of water, which resulted in an almost saturated solution. The sodium trithionate was then recrystallized by addition of ice cold absolute ethanol (at 0-2 °C). The volume of the ethanol added was approximately the same as that of the saturated solution. The purity of the crystallyzed sodium trithionate was checked by the following method. Known amount of sodium trithionate was dissolved in pure distilled water followed by addition of bromine solution in excess. The reaction mixture was left to stand for at least five minutes, which was proven to be enough to oxidize trithionate into sulfate completely. The excess of bromine was then removed by boiling the solution for a couple of minutes. The H+ liberated by the oxidation (S3O62- + 4Br2 + 6H2O f 3SO42- + 8Br- + 12H+) was determined by titrating the sample against standard sodium hydroxide solution. The validity of this stoichiometric equation was checked in separate experiments. The purity of the sodium trithionate was found to be better than 99.5%. Detailed investigations of the trithionate-bromine reaction are going to be carried out in our lab, and the results will be published later elsewhere. All the other chemicals (iodine, potassium iodide, acetic acid, sodium acetate, and sodium perchlorate) were of the highest purity commercially available and were used without further purification. The stock solutions were freshly prepared each day from double-distilled and twice ion-exchanged water.

10.1021/jp102661k  2010 American Chemical Society Published on Web 05/20/2010

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Figure 1. Measured (symbols) and calculated (solid lines) absorbance-time curves at pH ) 4.55, [S3O62-]0 ) 0.278 mM and [I2]0 ) 0.58 mM in the presence of initially added iodide ion. [I-]0/mM ) 0.0 (b), 0.1 (0), 0.3 ((), 0.7 (∆), 1.0 (9).

The pH of the solutions was regulated between 4.25 and 5.25 by acetic acid/acetate buffer taking the pKa of acetic acid as 4.55.15 The acetate concentration was always kept constant at 0.1 M and the pH was adjusted by the necessary amount of acetic acid. The ionic strength was set to 0.5 M by adding the necessary amount of sodium perchlorate. The temperature of the reaction vessel was maintained at 25.0 ( 0.1 °C. The initial concentrations of the reactants for all the 63 kinetic runs are collected in Table S1 of the Supporting Information. Methods and Instrumentation. The reaction was followed by a Zeiss S10 diode array spectrophotometer in the visible range without using the deuterium lamp. The reaction has been carried out in a standard quartz cuvette equipped with magnetic stirrer and Teflon cap having 1 cm optical path. The buffer components, the reactants iodine, and iodide (if necessary) were delivered from a pipet first. The spectrum of the solution was always recorded before injection of the trithionate solution to precisely determine the initial concentration of iodine. Then, the reaction was started with the addition of the necessary amount of trithionate solution from a fast delivery pipet. The spectra of the reacting solution at the wavelength range of 400-700 nm was acquired up to approximately 9000 s. To confirm the purity of the sodium trithionate and to check the sulfur-containing end products of the reaction, Raman spectroscopy was also carried out by a NXR FT-Raman spectrometer. Data Treatment. MRA studies have shown16 that the only absorbing species in the visible range are iodine and triiodide ion. Therefore, the simultaneous evaluation of the kinetic curves was carried out at the isosbestic point of I2-I3- system at 468 nm by the program package ZiTa.17 The molar absorbance of both species was found to be 750 M-1 cm-1 at this wavelength. Originally each kinetic run contained more than 400 absorbancetime data pairs, therefore it was necessary to reduce the number of time points (40-50) to avoid unnecessary time-consuming calculations. The essence of this method has already been described elsewhere.10 To obtain the kinetic model and the rate coefficients, absolute fitting procedure has been chosen to minimize the average deviation between the measured and calculated absorbance. Altogether, almost 2700 experimental points from the 63 kinetic series were used for the simultaneous evaluation. Our quantitative criterion for an acceptable fit was

that the average deviation for the absolute fit approached 0.004, which is close to the experimentally achievable limit of error. Results Stoichiometry. The stoichiometry of the reaction was established by standard iodometry as well as by Raman spectroscopy. Raman spectroscopy clearly showed that no sulfur-containing product formed via the reaction besides sulfate. After the reaction completed in excess of iodine the iodometric titration indicated that the strict stoichiometry of the reaction can be described by the following equation

S3O62- + 4I2 + 6H2O f 3SO42- + 8I- + 12H+

(1) This result was also confirmed by VIS-spectroscopy from the end of the kinetic measurements in excess of iodine because the consumed iodine-trithionate ratio calculated from the absorbance loss at 468 nm was found to be 4.0 ( 0.1. Since sulfate ion is the only sulfur containing end-product of the reaction that was supported by the Raman measurements (see: Figure S1 of the Supporting Information), eq 1 has to be also valid strictly in trithionate excess. Preliminary Observations. The most important characteristics and results of the preliminary individual curve fitting on the kinetic curves can be summarized as follows. (1) No pH dependence of the kinetic curves can be observed within the pH range studied (see: Figures S3 and S4 of the Supporting Information). This feature is not unique among the polythionate-iodine reactions because it was already noticed in the tetrathionate-iodine system as well.1,2 (2) Figure 1 shows the effect of initially added iodide ion on the decay of the total amount of iodine. It is easily seen that the increase of initial iodide concentration significantly retards the consumption of total amount of iodine. Our preliminary calculations demonstrated that eq 1, being treated as a simple second order chemical process (the kinetic order of both reactants is one) along with the fast equilibrium formation of triiodide ion, is not sufficient to interpret quantitatively the observed iodide inhibition. For the sake of completeness, the calculations yielded an unacceptably high 0.050 absorbance unit

Non-Triiodide Based Autoinhibition by Iodide Ion

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for the average deviation if the iodide inhibition is treated exclusively with the formation of triiodide. The nontriiodidebased inhibitory effect of iodide ion was already justified experimentally in the tetrathionate-iodine reaction as well.1,2 (3) The iodide inhibitory feature of the reaction made it possible to fit eq 1 with the following rate equation

TABLE 1: Fitted and Fixed Rate Coefficients of the Proposed Kinetic Modela step

rate equation

parameter

(R1) (-R1) (R2)

-

5.7 × 10 M-1 s-1 8.5 × 106 s-1 5.52 × 10-2 s-1 1.023 × 1011 M-2 s-1 1.98 × 10-3 Ms-1 3.67 × 109 M-1 s-1 103 M-1 s-1 107 M-1 s-1 24.7 ( 2.4 s-1 4.28 ( 0.51 M-1 s-1 (1.37 ( 0.09) × 10-3 s-1 (6.95 ( 1.01) × 103 M-1 s-1 3.1 × 109 M-1 s-1

kR1[I2][I ] k-R1[I3] kR2[I2] ′ kR2[HOI][H+][I-] k-R2[I2]/[H+] ′ k-R2[HOI][I-] kR3[S3O26 ][I2] k-R3[S3O6I-][I-] kR4[S3O6I-] k-R4[S3O6][I-] kR5[S3O6] kR6[SOI][I2] kR7[HSO3 ][I2]

(-R2)

V1 ) kapp

[S3O26 ][I2] [I-]

(2)

individually for most of the kinetic curves. The surprising result of our calculations was that with kapp ) (8.24 ( 1.50) × 10-4 s-1 all the measured 63 kinetic curves could be fitted with an almost acceptable average deviation of 0.0059 absorbance unit (see: Figures S5 and S6 of the Supporting Information). However, careful inspection of the individually fitted kinetic curves has revealed that systematic deviation could be noticed between the measured and calculated absorbance-time series, especially at large trithionate excess, high iodide concentrations, and relatively high reactant concentrations as well. More precisely, it meant that eq 2 was not able to describe almost one-third of the measured kinetic traces properly. We therefore concluded that eq 1 along with the rate equation indicated by eq 2 is not enough for sound description of the kinetic curves simultaneously, although it may serve as a solid starting point to create the final kinetic model. (4) In case of the remaining 20 kinetic curves (at higher initial iodine concentrations) we found that the following rate equation works adequately:

2-

V1 )

ka + kb[I-]

I2 + I- h I3

(R1)

I2 + H2O h HOI + I- + H+

(R2)

S3O26 + I2 h S3O6I + I

(R3)

S3O6I- h S3O6 + I-

(R4)

+ S3O6 + 3H2O f 3HSO3 + 3H

(R5)

2+ S3O6I- + I2 + 4H2O f 2HSO3 + SO4 + 3I + 6H (R6)

(R7)

(3)

with ka ) (1.51 ( 0.60) × 10-4 M2 s and kb ) (4.82 ( 1.74) × 10-2 M s. It should be emphasized that the second-order dependence of the rate equation on the concentration of iodine was found to be a necessary condition to fit those kinetic traces. Proposed Kinetic Model. Since the simple stoichiometric equation (eq 1) along with its rate equation (eq 2) was not able to fit the measured kinetic characteristics adequately, we have chosen the following set of species that most likely participate as a reactive intermediate. This group consists of S2O3OH-, S2O3I-, S3O6I-, S3O6, HSO3-, as well as HOI. Considering the possible mono- and bimolecular reactions among the reactants and the intermediates, we started the fitting procedure with 32 different chemically different steps. During the fitting procedure the steps, the rate coefficients of which became insensitive to the average deviation, were omitted step by step. This method has been applied successfully in several cases of our previous work.2,10,18 After long but straightforward systematic reduction, the following model has emerged.

CH3COOH h H+ + CH3COO-

a No error indicates that the value in question was fixed during the fitting procedure.

2+ HSO3 + I2 + H2O f SO4 + 2I + 3H

2

[S3O6 ][I2]

(R3) (-R3) (R4) (-R4) (R5) (R6) (R7)

9

(E1)

The rapid de- and reprotonation process E1 was taken into account with known equilibrium constant to follow the slight change in pH during the reaction. This acid dissociation equilibrium may be regarded as an auxiliary process necessary for the detailed calculation, but it is not central part of the proposed model. Table 1 contains the fitted and fixed kinetic parameters used in the simultaneous evaluation of the kinetic curves. The average deviation was found to be 0.0042 absorbance unit (AU). The final results of the proposed kinetic model are illustrated in Figures 1-4. Discussion Step R1 is the rapid equilibrium formation of triiodide ion investigated a couple of decades ago by several authors.19,20 The rate coefficients of the forward and reverse reactions were set to kR1 ) 5.7 × 109 M-1 s-1 and k-R1 ) 8.5 × 106 s-1, respectively to give log βI3- ) 2.83, where βI3- is the formation constant of triiodide ion.15 Step R2 is the well-known hydrolysis of iodine in aqueous solution. The rate coefficients of the forward and reverse reactions including the hydroxide-driven hydrolysis were determined previously, therefore we used these values as fixed parameters through the fitting procedure.21–23 Step R3 is the fast initiating equilibrium formation of S3O6Iexplained by a formal I+ transfer from iodine molecule to

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Figure 2. Measured (symbols) and calculated (solid lines) absorbance-time curves at pH ) 4.55 and [I2]0 ) 0.57 mM in absence of initially added iodide ion. [S3O62-]0/mM ) 0.155 (b), 0.10333 (0), 0.0775 (2), 0.062 ()), 0.05166 (9).

Figure 3. Measured (symbols) and calculated (solid lines) absorbance-time curves at [S3O62-]0 ) 0.085 mM, pH ) 4.55 and in absence of initially added iodide ion. [I2]0/mM ) 0.0666 (b), 0.148 (0), 0.308 (@), 0.473 ()), 0.643 (9), 0.72 (O).

possibly the inner sulfur atom of trithionate. This equilibrium is a complete analogy of the initiating equilibrium of the corresponding tetrathionate-iodine reaction.1,2 Our calculation indicated that this equilibrium must be shifted far to the left and the equilibrium constant cannot be determined from our experiments. If KR3 is set at any value less than 0.01, then exactly the same average deviation can be calculated, which means that we could set only an upper limit for KR3 with a sufficiently fast forward and backward reactions. To provide a sufficiently low level of [S3O6I-] during the calculation, we set KR3 ) 1 × 10-4 as a result of the ratio of kR3 and k-R3, being 103 M-1 s-1 and 107 M-1 s-1, respectively. Step R4 indicates that S3O6I- is involved in another equilibrium process in which cleavage of iodide ion results in a formation of an intermediate S3O6. The rate coefficients of the forward (kR4) and backward (k-R4) reactions can only be calculated if KR3 is fixed at 10-4 as mentioned above. In this case the rate coefficients of both the forward and reverse reaction could be determined from our experiments as 24.7 ( 2.4 s-1 and 4.28 ( 0.51 M-1 s-1, respectively. It should be emphasized

again that the value of kR4 is in total correlation with that of k-R3 as well as kR6, meaning that only their ratios (kR4/k-R3 and kR6/k-R3) could be calculated. Therefore, neither their absolute value nor the KR4 ) kR4/k-R4 equilibrium constant could be determined from our experiments. The underlying chemistry seems to be clear since it only means that S3O6I- must be a short-lived intermediate and its fast conversions govern the kinetics of the reaction. Further discussion about S3O6 will be detailed later in the formal kinetics part. Step R5 is the hydrolysis of S3O6 formed in step R4 and kR5 was determined to be (1.37 ( 0.09) × 10-3 s-1. This pathway is responsible for governing the reaction at lower iodine and initial iodide concentrations, that is, at larger trithionate excess. It is also interesting to note that steps R3 and R4 may be equilibrated in the initial phase of the reaction in such a way that the concentration of S3O6 might be as high as 10-4 M, since its lone pathway to decay its concentration is relatively slow. The halftime of S3O6 calculated from kR5 is about 8 min meaning that it may be accumulated in some extent during the first stage of the reaction. Therefore, we tried to exchange this reaction

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Figure 4. Measured (symbols) and calculated (solid lines) absorbance-time curves at constant [S3O62-]0/[I2]0 ≈ 1.0 ratio in absence of initially added iodide ion at pH ) 4.55. [I2]0/mM ) 0.0737 (O), 0.157 (×), 0.238 (∆), 0.328 (+), 0.518 (b), 0.588 (0), 0.635 (2), 0.68 ()), 0.756 (9).

by many conceivable reaction steps with no success. Exclusion of this reaction results in not even an unacceptably high average deviation (0.0209 AU) but also systematic deviations between the measured and calculated absorbance-time curves especially at high initial trithionate concentrations. If, however, kR5 is supposed to be a fast reaction (kR5 > 102 s-1) then the average deviation was found to be 0.0081 AU, but systematic deviation can still be observed between the measured and calculated kinetic curves (see: Figure S7 of the Supporting Information). These facts together serve as an indirect evidence for the necessity of the inclusion of step R5, therefore we concluded this step to be necessary for detailed and sound description of the kinetic curves. Step R6 is a process of fast oxidation of S3O6I- by iodine. As mentioned above, the absolute value of this rate coefficient cannot be determined from our experiments, only the ratio of kR6/k-R3. Since k-R3 was fixed at 107 M-1 s-1, our calculation has provided a value of (6.95 ( 1.01) × 103 M-1 s-1 for kR6. This pathway determines the kinetics at higher initial iodine concentrations. It is interesting to note that we have also tried to replace step R6 by the following equation:

S3O6I- + 3H2O f 3HSO3- + I- + 3H+

(4)

This change, however, would result in a relatively small increase in the average deviation from 0.0042 to 0.0054 absorbance unit, but at the highest initial iodine concentrations systematic deviations could be encountered between the measured and calculated kinetic curves. Thus, it is concluded that Step R6 cannot be replaced by eq 4. Step R7 is the well-known fast reaction between hydrogen sulfite and iodine. This reaction was thoroughly studied by Yiin and Margerum,24 and even the rate coefficient of this process was determined to be 3.1 × 109 M-1 s-1. Therefore, we have fixed kR18 at this value. Formal Kinetics. As one may easily notice S3O6I- and HSO-3 are short-lived intermediates, therefore steady-state approximation may be applied with no restrictions to derive a formal rate equation. However, for S3O6 the steady-state approximation is held only at high initial iodide concentrations or at low initial

trithionate and iodine concentrations. At such experimental conditions the following rate law can be deduced:

-

{

}

kR6 kR6 1 d[I2] [I2][I-] + [I ] / ) kR3 1 + 4 dt kR5KR4 kR4 2 kR6 kR6 [I2][I-] + [I ] + 1+ kR5KR4 kR4 2 k-R3 - 2 k-R3 [I ] + [I ] [S3O62-][I2] kR4 kR5KR4

{

}

(5)

The second and the third terms in the nominator of the righthand side of eq 5 is negligible because the conditions of [I2] , (kR4)/(kR6) ) 0.00355 M and [I2][I-] , (KR5kR4)/(kR6) ) 1.14 × 10-6 M2 are fulfilled in most of the kinetic traces measured. Moreover [I-] . (kR4)/(k-R3) ) 2.47 × 10-6 M is also valid as a result of the relatively rapidly established iodine hydrolysis represented by eq R2. It allows us to simplify eq 5 as

-

d[S3O62-] 1 d[I2] )) dt 4 dt kR3

1 [S O 2-][I2] k-R3 - 2 3 6 k-R3 [I ] + [I ] kR4 kR5KR4

(6)

since the first three terms in the denominator of eq 5 are also negligible. After some algebraic manipulation the following equation can be obtained:

-

[S3O62-][I2] d[S3O2kR5kR4 6 ] 1 d[I2] )) KR3 dt 4 dt kR5 + k-R4[I-] [I-] (7)

Comparing eq 2 to eq 7 reveals that

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kapp )

KR3kR5kR4 -

kR5 + k-R4[I ]

Cseko˝ and Horva´th

(8)

that is, the apparent rate coefficient obtained by the individual curve fitting method must depend on the actual concentration of iodide ion. This feature of the apparent rate coefficient is also well-established in the corresponding tetrathionate-iodine reaction.1,2 Substitution of the corresponding rate coefficients into eq 8 and taking into consideration that the average iodide concentration during most of the kinetic runs is approximately 6 × 10-4 M, one can calculate 8.59 × 10-4 s-1 for kapp, which is in complete coincidence with the fact that most of the kinetic runs may be evaluated by eq 2 found in the preliminary studies. This expression also provides a smooth understanding why the individual rate coefficient of equilibrium R3 cannot be determined from these experiments. If the value of KR3 is fixed at a relatively small value to maintain the concentration of S3O6Iat a sufficiently low level than the values of kR4, k-R4, and kR5 can straightforwardly be determined. To check it, we have undertaken an additional simultaneous curve fitting by the proposed model not including step R6 only on those kinetic curves, where the initial iodine concentration was lower than 6 × 10-4 M. This condition meant that only 43 kinetic curves were used for parameter estimation. As expected, the remaining three parameters are also capable of good description of the selected kinetic curves with almost the same average deviation (0.0043 AU). It is therefore concluded that under such experimental conditions the value of kR6 cannot be calculated. The underlying chemistry seems to be clear since at low iodine concentrations the main pathway is Step R5 to produce sulfate and iodide eventually. It, however, raises an important question about the determination of kR6 by nonlinear parameter estimation under present circumstances. As already mentioned previously, not all the kinetic runs can be evaluated properly by eq 2. It is easily seen that the second term of the nominator as well as that of the denominator of eq 5 becomes no longer negligible with an increase of the initial iodine concentration. It straightforwardly introduces the dependence of kapp on kR6, which provides a solid base to determine the value of kR6 under our experimental conditions by the remaining 20 kinetic curves having high initial iodine concentrations. Thus, we concluded that the 63 kinetic curves carry enough experimental information to evaluate the parameters kR4, k-R4, kR5, and kR6 if KR3 is fixed. Finally, it should also be noted that the formal rate equation derived above is not valid at larger reactant concentrations because one of the intermediates may be accumulated (S3O6) in an appreciable amount in the early stage of the reaction. Therefore, proper explanation of the necessity of the secondorder dependence of iodine in the rate equation is not possible via simplification of eq 5, although mathematically it can easily be deduced if the first term of the nominator in eq 5 is neglected compared to the terms depending on the concentration of iodine. Conclusion The work presented here may be considered as the first reliable attempt to unravel the kinetics and mechanism of the trithionate-iodine reaction. It was experimentally demonstrated that the reaction is autoinhibitory with respect to iodide ion, and this effect cannot be solely attributed to the formation of triiodide ion, which is less reactive toward the substrate than iodine. The nontriiodide-based inhibition of iodide ion is not a

unique feature; it rather seems to be a general phenomenon among the oxidation of polythionates by iodine that is based on a formal I+ transfer from iodine molecule to (one of) the inner sulfur of polytionates as an initiating equilibrium that was previously established in the tetrathionate-iodine reaction as well.1,2 The continuation of elucidation of this type of reaction provides a better understanding and insight into the nature of the sulfur-chain breaking reactions of polythionates that may also be helpful as a guide in unraveling the more complex oxidation reactions of these compounds. A better insight into the mechanism of these reactions would thus result in an easier and model-based comprehensive explanation of the wide variety of the nonlinear dynamical phenomena appearing in a reaction4–6,12 having rather complex kinetics. Acknowledgment. This work was supported by the Hungarian Research Fund (OTKA) Grant No. K68172. A. K. H. is grateful for the financial support of the Ja´nos Bolyai Research Scholarship of the Hungarian Academy of Sciences. Supporting Information Available: Table containing the condition of each kinetic run and figures containing the measured and fitted data are found in the Supporting Information. This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes (1) Awtrey, A. D.; Connick, R. E. J. Am. Chem. Soc. 1951, 73, 4546. (2) Kerek, A.; Horváth, A. K. J. Phys. Chem. A 2007, 111, 4235– 4241. (3) Horva´th, A. K.; Nagypa´l, I.; Peintler, G.; Epstein, I. R. J. Am. Chem. Soc. 2004, 126, 6246–6247. (4) Horva´th, D.; Tóth, A. J. Chem. Phys. 1997, 108, 1447–1451. (5) Fuentes, M.; Kuperman, M. N.; Kepper, P. D. J. Phys. Chem A 2001, 105, 6769–6774. (6) Lima, D.; D’Onofrio, A.; Wit, A. D. J. Chem. Phys. 2006, 124, 014509. (7) Boissonade, J.; Dulos, E.; Gauffre, F.; Kuperman, M. N.; Kepper, P. D. Faraday Discuss. 2002, 120, 353–361. (8) Szalai, I.; Gauffre, F.; Labrot, V.; Boissonade, J.; Kepper, P. D. J. Phys. Chem. A 2005, 109, 7843–7849. (9) Horva´th, A. K. J. Phys. Chem. A 2005, 109, 5124–5128. (10) Horva´th, A. K.; Nagypa´l, I.; Epstein, I. R. Inorg. Chem. 2006, 45, 9877–9883. (11) To´th, A.; Horva´th, D.; Siska, A. J. Chem. Soc. Faraday Trans. 1997, 93, 73–76. (12) Peintler, G.; Cseko˝, G.; Petz, A.; Horváth, A. K. Phys. Chem. Chem. Phys. 2010, 12, 2356–2364. (13) (a) Gutman, A. Ber. Dtsch. Chem. Ges. 1906, 39, 509. (b) Riesenfeld, E. H. Z. Anorg. Allg. Chem. 1924, 141, 109. (14) Kelly, D. P.; Wood, P. A. Methods Enzymol. 1994, 234, 475–501. (15) IUPAC Stability Constant Database; Royal Society of Chemistry: London, 1992-1997. (16) Peintler, G.; Nagypa´l, I.; Jancso´, A.; Epstein, I. R.; Kustin, K. J. Phys. Chem. A 1997, 103, 8013–8020. (17) Peintler, G., ZiTa, Version 5.0: A ComprehensiVe Program Package for Fitting Parameters of Chemical Reaction Mechanism; Attila Jo´zsef University: Szeged, Hungary, 1989-1998. (18) (a) Varga, D.; Horva´th, A. K. J. Phys. Chem. A 2009, 113, 9988– 9996. (b) Varga, D.; Horváth, A. K. J. Phys. Chem. A 2009, 113, 13907– 13912. (19) Turner, D. H.; Flynn, G. W.; Sutin, N.; Beitz, J. V. J. Am. Chem. Soc. 1972, 94, 1554–1559. (20) Ruasse, M.; Aubard, J.; Galland, B.; Adenier, A. J. Phys. Chem. 1986, 90, 4382–4388. (21) Eigen, M.; Kustin, K. J. Am. Chem. Soc. 1962, 84, 1355–1361. (22) Lengyel, I.; Epstein, I. R.; Kustin, K. Inorg. Chem. 1993, 32, 5880– 5882. (23) Lengyel, I.; Li, J.; Kustin, K.; Epstein, I. R. J. Am. Chem. Soc. 1996, 118, 3708–3719. (24) Yiin, B. S.; Margerum, D. W. Inorg. Chem. 1990, 29, 1559–1564.

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