Kinetics and Mechanism of the Oxidation of Bromide by Periodate in

Article pubs.acs.org/JPCA

Kinetics and Mechanism of the Oxidation of Bromide by Periodate in Aqueous Acidic Solution Viktor Szél, György Csekő, and Attila K. Horváth* Department of Inorganic Chemistry, University of Pécs, Ifjúság útja 6, H-7624 Pécs, Hungary S Supporting Information *

ABSTRACT: The periodate−bromide reaction has been studied spectrophotometrically mainly in excess of bromide ion, monitoring the formation of the total amount of bromine at 450 nm at acidic buffered conditions and at a constant ionic strength in the presence of a phosphoric acid/dihydrogen phosphate buffer. The stoichiometry of the reaction was established to be strictly IO−4 + 2Br− + 2H+ → Br2 + IO−3 + H2O. The formal kinetic order of the reactants was found to be perfectly one and two in the cases of periodate and bromide, respectively, but that of the hydrogen ion lies between one and two. We have also provided experimental evidence that dihydrogen phosphate accelerates the formation of bromine, suggesting the appearance of strong buffer assistance. On the basis of the experiments, a simple two-step kinetic model is proposed involving BrIO3 as a key intermediate that perfectly explains all of the experimental findings. Furthermore, we have also shown that in huge excess of bromide, the apparent rate coefficient obtained from the individual curve fitting method of the absorbance−time series is necessarily independent of the initial periodate concentration that may falsely be interpreted as the rate of bromine formation is also independent of the concentration of periodate.

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INTRODUCTION Periodate is a powerful oxidizing agent in dilute aqueous solutions, and therefore, it has long been the subject of studies of elaborating different analytical procedures.1,2 The majority of these methods is based on the direct oxidation of iodide and bromide ions by periodate producing elementary halogens that interact with the species to be analyzed quantitatively by different methods. For example, it has been shown that trace amounts of arsenic,3 thiocyanate,4 and phenylhydrazine5 and micro amounts of bromide and iodide ions6 can conveniently and precisely be determined by these methods. Furthermore, there has been a growing interest in the application of the bromide/periodate mixture in organic synthesis due to in situ generation of bromine under mild conditions.7−11 Despite these widespread interests in the application of the bromide/ periodate mixture in analytical and organic chemistry, no direct detailed kinetic study of the periodate−bromide reaction is yet available in the literature. Compared to that, the kinetics of the corresponding iodide−periodate reaction has long been disputed with regard to the H+ dependence of its rate.12−14 Indelli et al. came to a conclusion that the rate of the periodate−iodide reaction is linearly proportional to [H+] based on their initial rate studies,12,14 but Marques and Hasty’s result13 was in perfect harmony with the early measurements15 of Abel and Siebenschein indicating lack of pH dependence in slightly acidic medium. Later, it turned out that the rate of formation of iodine is independent of pH, but the stoichiometry of the reaction strongly depends on the pH at slightly acidic conditions.16 The disagreement arose from the fact that the amount of iodine formed in a time unit at the © 2014 American Chemical Society

initial phase of the reaction increased with decreasing pH due to the shift of the actual stoichiometry of the reaction to the limiting stoichiometry producing more iodine. If anyone falsely supposes a strict stoichiometry to the reaction, it would be misinterpreted that the rate of the reaction depends on pH. Lack of a detailed study of the title reaction and the history of the dispute on the corresponding iodide−periodate reaction prompted us to elucidate the kinetics and mechanism of the title reaction. Furthermore, especially in the case of applying the reaction for bromination of organic substances, it is believed that elementary bromine and iodide ions form in the initiating reaction,9 which is difficult to accept in view of the fact that bromine, being a very strong oxidizing agent, would react rapidly with the iodide ion to give molecular iodine.17

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EXPERIMENTAL SECTION Materials and Buffers. All of the reagents (sodium bromide, potassium metaperiodate, sodium dihydrogen phosphate, phosphoric acid, potassium bromate, potassium bromide, and sodium perchlorate) were of the highest purity available commercially and were used without further purification. Stock solutions were prepared from double-ion-exchanged and double-distilled water. The pH of the solutions was regulated between 1.5 and 2.3, taking the pKa of phosphoric acid to be 1.8.18 The ionic strength was controlled at 0.5 M at each experiment by addition of the necessary amount of sodium Received: September 10, 2014 Revised: October 21, 2014 Published: October 22, 2014 10713

dx.doi.org/10.1021/jp509164e | J. Phys. Chem. A 2014, 118, 10713−10719

The Journal of Physical Chemistry A

Article

IO−4 + 2Br − + 2H+ = IO−3 + Br2 + H 2O

perchlorate. The temperature of the reaction vessel was kept constant at 25.0 ± 0.1 °C. The initial concentration of the reactants was varied between 0.6 and 6.0 mM and 7.8 and 110 mM in the cases periodate and bromide, respectively. The stock bromide and periodate solutions were prepared by weighing the necessary amounts of solid substances. Because the majority of the measurements was carried out in an excess of bromide, it was also necessary to determine the stability constant of the tribromide ion. The bromine− tribromide system was in situ prepared by acidifying the mixture containing the necessary amount of potassium bromate and potassium bromide (at 11 different initial conditions; see the Supporting Information), and the samples were left to stand for 3 days to complete the reaction. Methods and Instruments. The reaction was followed by a Zeiss S600 diode array spectrophotometer in the visible range without using the deuterium lamp of the instrument. A quartz cuvette, equipped with a Teflon cap having a 1 cm path length, was used as a reactor. The solution was continuously stirred at 750 rpm by a small magnetic bar. The buffer solution was first delivered into the cuvette followed by the necessary amount of the reactant to be held constant during the series of the measurements. Finally, the reaction was started by injecting the necessary amount of the other reactant from a fast delivery pipet. We have also checked the reproducibility of the kinetic curves in several different cases. These measurements convinced us about a very good reproducibility (see Figure 7). The spectrum of the reaction solutions at the wavelength range of 400−800 nm was acquired up to approximately 800− 240000 s depending on the initial conditions. Data Treatment. In the visible range, only bromine and tribromide were found to be the absorbing species; therefore, the isosbestic point of the bromine−tribromide system (λ = 450 nm) was selected for the parameter estimation by the ChemMech/ZiTa program package developed to fit a basically unlimited experimental series.19 Because the rigorous fitting procedure requires precise molar absorbance and stability constant values, the bromine−tribromide equilibrium system was evaluated by the PSEQUAD program package.20 The result of the fit can be seen in the Supporting Information. We obtained for the molar absorbance of the bromine and tribromide ion a 104 ± 1 M−1cm−1 value, and for βBr−3 , we obtained a value of 18.06 ± 0.06 M−1, which were found to be in a very good agreement with the values obtained from independent studies.21,22 Originally, each kinetic run contained more than 500 absorbance−time data pairs; therefore, it was necessary to reduce the number of time points (40−80) to avoid unnecessary and time-consuming calculations. The essence of this method has already been described elsewhere.23 Altogether, almost 3724 experimental points form the 70 kinetic series were used for the simultaneous evaluation. Our quantitative criterion for an acceptable fit was that the average deviation for the absolute fit approach 0.004 absorbance unit, which is close to the experimentally achievable limit of error of the spectrophotometer.

(1)

Our kinetic measurements also strengthen this fact, as can be seen in Table 1, because from the end of the spectrophotoTable 1. Calculated Stoichiometric Ratio (SR) [IO−4 ]0/mM

[Br−]0/mM

pH

TBr2/mM

SR

0.624 0.936 5.81 3.12 3.21 1.56 2.18 4.37 3.12 3.12 0.624 2.18

78.2 78.2 78.2 15.6 109.5 78.2 78.2 78.2 31.2 78.2 78.2 78.2

1.52 1.52 1.52 1.70 1.70 1.89 1.89 1.89 2.08 2.08 2.28 2.28

0.625 0.947 5.885 3.07 3.16 1.51 2.12 4.33 3.03 3.10 0.620 2.13

1.00 1.01 1.01 0.98 0.98 0.97 0.97 0.99 0.97 0.99 0.99 0.98

metric measurement, the final concentration of total bromine (TBr2 = [Br2] + [Br3−]) can easily be calculated. The stoichiometric ratio (SR), defined as TBr2/[IO4−]0, shows unity, perfectly supporting eq 1 as the stoichiometry of the reaction. Initial Rate Studies. Figures 1 and 2 indicate the result of the initial rate studies. The calculated initial rates were obtained

Figure 1. Initial rate studies to determine the formal kinetic order of periodate at five different pHs. Conditions are as follows: [Br−]0 = 78.2 mM; pH = 1.52 (black), 1.70 (blue), 1.89 (green), 2.08 (cyan), and 2.28 (red).

from the slope of the measured absorbances at 450 nm extrapolated to t = 0 s. It is clear that the formal kinetic orders of periodate and bromide ions are strictly one and two, respectively. The formal kinetic order of the hydrogen ion was, however, found to be between one and two (see Figure 3), suggesting a complex mechanism with respect to this species. Later, we shall see that this experimental finding can easily be explained by our proposed kinetic model. Individual Curve Fitting Method. Because the majority of the kinetic curves was measured in a high excess of bromide, it provided an opportunity to fit each absorbance−time profile with the following equation

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RESULTS Stoichiometry. It is already well-established that bromine (and tribromide ion) and iodate form as products of the reaction; therefore, the stoichiometry can strictly be characterized by the following equation6,24,25 10714



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Figure 4. Logarithm of the apparent rate coefficient against the logarithm of reactant concentrations. Conditions are as follows: pH = 1.70, [Br−]0 = 78.2 mM (blue), [IO−4 ]0 = 3.12 mM, pH = 1.70 (black); [Br−]0 = 78.2 mM, [IO−4 ]0 = 2.18 mM (green).

Figure 2. Initial rate studies to determine the formal kinetic order of bromide at four different pHs. Conditions are as follows: [IO−4 ]0 = 3.12 mM; pH = 1.52 (black), 1.70 (blue), 1.89 (green), and 2.08 (red).

Figure 3. Initial rate studies to determine the formal kinetic order of the hydrogen ion. Conditions are as follows: [IO−4 ]0 = 2.18 mM, [Br−]0 = 78.2 mM (black); [IO−4 ]0 = 3.12 mM, [Br−]0 = 31.3 mM (blue).

A = A 0(1 − e−kappt )

Figure 5. Logarithm of the initial rate against the logarithm of buffer concentrations. Conditions are as follows: pH = 2.35, [Br−]0 = 78.2 mM, [IO−4 ]0 = 0.63 mM, I = 0.5 M. The slope was found to be 0.385 ± 0.018.

(2)

consideration because the rate of the formation of bromine strongly increases with increasing buffer concentration. Proposed Kinetic Model. To propose the kinetic model, we suggest a rapidly established preequilibrium between the reactants producing a short-lived intermediate BrIO3. This species is analogous of I2O2 proposed to form in the initiating equilibrium in the Dushman reaction.26 It should also be emphasized that the effect of phosphate buffer catalysis27 is suggested to be taken into consideration by the subsequent transition of the I2O2 intermediate in the case of the iodide− iodate reaction.26 Therefore, as a conceivable analogy, we propose the formation of BrIO3 in an initiating equilibrium, the species that further reacts with another bromide ion to produce bromine and iodate via possibly a bromonium ion transfer process in a subsequent reaction. We also suggest that this reaction is accelerated by the buffer component H2PO−4 , like in the case of the analogous Dushman reaction.

Table S1 of the Supporting Information summarizes the results of the individual fit. It clearly indicates that the apparent rate coefficients are proportional to the bromide and hydrogen ion concentrations, but it suggests that the apparent rate coefficient is independent of the initial periodate concentration at constant bromide and hydrogen ion concentrations. Figure 4 further visualizes the effect of the initial concentration of the reactants on the apparent rate coefficients obtained. As can be seen, comparing the slopes of the fitted straight lines of the log−log plots with the ones obtained from the initial rate studies reveals that the formal kinetic orders of hydrogen and bromide ions are found to be in perfect harmony with each other in these cases. An apparent contradiction may, however, be recognized between the unity of the formal kinetic order of the periodate ion obtained from the initial rate studies and the periodate independence of kapp. Later, we shall see that these findings can straightforwardly be reconciled with the proposed kinetic model. Effect of Buffer Concentration. Figure 5 indicates the dependence of the initial rate with respect to the concentration of dihydrogen phosphate, keeping the pH, the reactant concentrations, and the ionic strength constant. It clearly indicates that a strong buffer assistance has to be taken into

H3PO4 ⇌ H+ + H 2PO−4

(E1)

Br − + IO−4 + 2H+ ⇌ BrIO3 + H 2O

(R1)

−

BrIO3 + Br → Br2 +

IO−3

(R2)

Rate coefficients determined by nonlinear simultaneous parameter estimation are illustrated in Table 2. 10715



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Table 2. Fitted and Fixed Rate Coefficients of the Proposed Kinetic Modela step

rate equation

parameter value

R1

kR1[IO−4 ][Br−][H+]2 k−R1[BrIO3] kR2[BrIO3][Br−] kR2 ′ [BrIO3][Br−][H2PO−4 ][H+]−1

107 M−3 s−1 109 s−1 (1.46 ± 0.13) × 105 M−1 s−1 947 ± 80 M−1 s−1

R2 R2′

No error indicates that the given parameter is fixed during the calculation process. a

The average deviation was found to be 0.0035 absorbance units by an absolute fitting procedure. Altogether, only two fitted parameters were used. Figures 6−9 demonstrate the quality of the fit for representative examples and also support the fact that the proposed kinetic model is working properly under our experimental conditions.

Figure 8. Experimental (symbols) and calculated (solid lines) absorbance−time curves. Conditions are as follows: [IO−4 ]0 = 4.37 mM, [Br−]0 = 78.2 mM, pH = 1.52 (black), 1.70 (blue), 1.89 (green), 2.08 (cyan), and 2.28 (red); I = 0.5 M.

Figure 6. Experimental (symbols) and calculated (solid lines) absorbance−time curves. Conditions are as follows: pH = 1.50, [IO−4 ]0 = 3.12 mM, [Br−]0/mM = 7.82 (black), 10.9 (blue), 15.6 (green), 21.9 (cyan), 31.2 (red), 46.9 (magenta), 78.2 (brown), and 109.5 (yellow); I = 0.5 M.

Figure 9. Experimental (symbols) and calculated (solid lines) absorbance−time curves. Effect of the buffer concentration. Conditions are as follows: pH = 2.35, [Br−]0 = 78.2 mM, [IO−4 ]0 = 0.63 mM, I = 0.5 M. [H2PO−4 ]0/mM = 2.81 (black), 11.2 (blue), 22.5 (green), 45.0 (cyan), and 90.0 (red).

[H 2PO−4 ] ⎞ d[Br2] ⎛ − ′ = ⎜kR2 + kR2 ⎟[BrIO3][Br ] ⎝ dt [H+] ⎠

(3)

Applying steady-state approximation for the short-lived BrIO3 intermediate and bearing in mind that k−R1 ≫ (kR2 + kR2 ′ [H2PO−4 ]/[H+]) under our experimental conditions leads to the expression of [BrIO3] = KR1[Br −][IO−4 ][H+]2

(4)

where KR1 = kR1/k−R1. Substituting eq 4 into eq 3 leads to [H 2PO−4 ] ⎞ d[Br2] ⎛ − 2 − + 2 ′ = ⎜kR2 + kR2 ⎟KR1[Br ] [IO4 ][H ] ⎝ dt [H+] ⎠

Figure 7. Experimental (symbols) and calculated (solid lines) absorbance−time curves. Conditions are as follows: pH = 1.89, [Br−]0 = 78.2 mM, [IO−4 ]0/mM = 0.624 (black), 0.936 (blue), 1.56 (green), 2.18 (cyan), 3.12 (red,yellow), 4.37 (magenta), and 5.93 (brown); I = 0.5 M.

(5)

Equation 5 clearly explains the formal kinetic orders of the bromide ion and periodate ion to be two and one, respectively, as well as the kinetic order of the hydrogen ion to be between one and two. It also shows that the formation of bromine is affected by buffer catalysis. Moreover, it is clear that in fact only KR1kR2 = 1460 ± 130 M−4 s−1 and KR1k′R2 = 9.47 ± 0.80 M−4 s−1 can be calculated from our measurements. As expected, KR1 has to be chosen for an arbitrarily small value to fulfill the

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DISCUSSION First, we shall show that the proposed model explains all of the experimental findings obtained from the initial rate studies. Formation of bromine can be expressed as 10716



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equilibrium (step R1) occurs in the following sequence of equations and thus accounts also for the second-order dependence of the hydrogen ion:

criterion of applying low steady-state concentration for species BrIO3. Second, we shall also show that kapp determined from the individual curve fitting method should necessarily be independent of the initial concentration of the periodate ion. For a C → B first-order reaction, analytical solution of the differential equation leads to the following expression [B] = [C]0 (1 − e−kappt )

KD

IO−4 + 2H 2O ⇌ H4IO−6 K1

H5IO6 ⇌ H4IO−6 + H+

(6)

(9)

K10

H5IO6 + Br − + H+ HooI BrIO3 + 3H 2O

Converting the concentrations into absorbance leads to the following expression A = εl[B] = εl[C]0 (1 − e−kappt )

(8)

(10)

where K10 = KR1K1/KD. Consequently, the kinetically active species is supposed to be rather the form of H5IO6 than the metaperiodate ion. A word is also in order here to enlighten the similarities and differences between the iodide−periodate and bromide−periodate reactions. It was previously shown that the iodide−periodate reaction proceeds via a pH-independent pathway at slightly acidic conditions.16 In contrast to this, the present reaction was found to be vanishingly slow under this condition, meaning an essential difference between these reactions. However, at strongly acidic conditions, the rate of the iodide−periodate reaction was also found to be proportional15 to [H+]2, the same as was found in the case of the title reaction. It appears to suggest that at strongly acidic conditions in case of both reactions, orthoperiodic acid is the kinetically active species, while at slightly acidic conditions, the metaperiodate ion seems to react with the iodide ion but not with the bromide ion. Certainly, at this point, it is only a reasonable explanation that further experiments are required to confirm this idea. Finally, it should also be mentioned that the proposed kinetic model is not the only possibility to describe the kinetic feature of the reaction. Proposing a short-lived intermediate BrIO3 is only a single feasibility. Here, we, however, provide two reasonable arguments why this species may be a reasonable candidate as a key intermediate. First, involving Cl2O3 as a short-live transient species in the decomposition of chlorous acid is well-established.29 Second, the strong buffer assistance observed, however, makes this possibility quite likely as well because subsequent reaction of similar intermediates, like Cl2O2, Br2O2, and ClBrO2, I2O2, is frequently accompanied by buffer catalysis.26,30−32 Such analogies thus would seem to favor the inclusion of BrIO3 as a key intermediate in the title reaction. Because the bromonium ion (Br+) is expected to be a good leaving group from BrIO3, we suggest that this species may partially be responsible for bromination of the several organic substances at the molecular level.

(7)

Because the path length of the cuvette is l = 1 cm, plotting the fitted A0 (obtained from the individual curve fitting method based on eq 2) against the initial concentration of periodate should give a perfectly straight line starting from zero and having the slope equal to the molar absorbance of bromine (or tribromide) at the isosbestic point. As a result, we prepared Figure 10, indicating the A0 versus [IO−4 ]0 plot from our

Figure 10. Plot of A0 determined from the individual curve fitting method against the initial concentration of the periodate ion. Conditions are as follows: [Br−]0 = 78.2 mM, pH = 1.52.

measurements that perfectly supports this fact if one compares the slope of the fitted straight line 104.6 M−1 cm−1 with εBr2 = εBr3− = 104 M−1 cm−1 obtained from the equilibrium measurements independently. Furthermore, it also indicates two important consequences: first, (a) in an excess of bromide ion, the total amount of bromine (TBr2) is determined by the initial concentration of periodate during the whole course of the reaction; consequently, no intermediate accumulates during the course of the reaction in detectable amount; and second, (b) the kinetic order of periodate being unity, determined from the initial rate study, should necessarily mean that in an excess of bromide ion, the kapp obtained from the individual curve fitting method is independent of [IO−4 ]0. kapp depends only on the pH and the concentration of bromide ion. Of course, if the huge bromide ion excess is not maintained anymore, eq 7 is no longer valid due to the fact that the pseudo-first-order conditions are not fulfilled; consequently, the individual fit of the kinetic curves cannot be performed by eq 7. It is also interesting to note that the periodate ion in acidic aqueous solution under our experimental condition mainly occurs in the form of H5IO6.28 It follows from the fact the pK1 and pK2 of orthoperiodic acid are 3.3 and 6.7, respectively.28 Having this information in hand, we suggest that the initiating

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CONCLUSION Here, we have reported the results of the detailed kinetic investigation of the bromide−periodate reaction in acidic buffered conditions. The reaction is governed by a strict stoichiometry producing bromine and iodate ion, and the reaction is first- and second-order with respect to the periodate and bromide ions. The rate of the reaction strongly depends on pH as well as on the concentration of buffer components, suggesting a strong buffer assistance. Besides providing the most important kinetic information on this practically important reaction, an additional impact seems to be worth highlighting. This paper clearly shows that implications drawn from the so-called simplified evaluation techniques in chemical kinetics must be used with special circumspection even in the case of the simplest reaction. The seemingly contradictory results may easily be reconciled by a correct interpretation and 10717



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understanding of the experimental findings that may easily help to avoid coming up unintentionally with sloppy conclusions. As an example, we would like to refer to the main conclusion of Christy and Egeberg’s study,33 who concluded that the thiocyanate−hydrogen peroxide reaction (studied under high excess of hydrogen peroxide) was zeroth-order with respect to thiocyanate,33 contradicting all of the earlier and recent studies on this subject.34−36 Inspecting their paper seems to be very reminiscent of the main outcome that we presented here because the apparent rate coefficients were also determined by an individual curve fitting method in that case. Therefore, it is highly important to emphasize that in a huge excess of a reactant, the apparent rate coefficients, determined from the concentration−time series of the limiting agent or that of a product having its concentration directly proportional to the limiting agent, obtained by an individual curve fitting method, necessarily have to be independent of the concentration of the limiting species (of course, only if high excess of the other reactant is maintained), but it does not mean that the reaction is zeroth-order with respect to the limiting agent.

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(10) Mohammad, A. Z.; Mojtaba, B.; Shadpour, M.; Gholamabbas, C.; Arash, G.; Nadiya, K.; Mina, D.; Mahmoudreza, D. The First Report on the Catalytic Oxidation of Urazoles to Their Corresponding Triazolinediones via In Situ Catalytic Generation of Br+ Using Periodic Acid or Oxone/KBr System. J. Mol. Catal. A: Chem. 2007, 270, 219− 224. (11) Kumar, L.; Mahajan, T.; Agarwal, D. D. Bromination of Deactivated Aromatic Compounds with Sodium Bromide/Sodium Periodate under Mild Acidic Conditions. Ind. Eng. Chem. Res. 2012, 51, 11593−11597. (12) Indelli, A.; Ferranti, F.; Secco, F. A Kinetic Study of the Reaction of Periodate with Iodide Ions. J. Phys. Chem. 1966, 70, 631− 636. (13) Marques, C.; Hasty, R. A. Application of the Iodide-IonSelective Electrode to a Kinetic Study of the Periodate−Iodide Reaction. J. Chem. Soc., Dalton Trans. 1980, 1269−1271. (14) Ferranti, F.; Indelli, A. On the Kinetics of the Periodate + Iodide Reaction. J. Chem. Soc., Dalton Trans. 1984, 1773−1774. (15) Abel, E.; Siebenschein, R. Ermittlung Zeitlich Unzuganglicher Reaktionkinetik Durch Reaktionsverteilung. Z. Phys. Chem. 1927, 130, 631−657. (16) Horváth, A. K. Pitfall of an Initial Rate Study: On the Kinetics and Mechanism of the Reaction of Periodate with Iodide Ions in a Slightly Acidic Medium. J. Phys. Chem. A 2007, 111, 890−896. (17) Rao, T. S.; Mali, S. I.; Dangat, V. T. Kinetics of the Rapid Reaction Br2 + 2I− = I2 + 2Br− in Aqueous Solution. Z. Phys. Chem. 1979, 260, 38−42. (18) IUPAC Stability Constant Database; Royal Society of Chemistry: London, 1992−1997. (19) Peintler, G. ChemMech/ZiTa; A Comprehensive Program Package for Fitting Parameters of Chemical Reaction Mechanism; Attila József University: Szeged, Hungary, 1989−2011. (20) Zékány, L.; Nagypál, I.; Peintler, G. PSEQUAD for Chemical Equilibria; Technical Software Distributors: Hungary, 1991. (21) Wang, T.; Kelley, M. D.; Cooper, J. N.; Beckwidth, R. C.; Margerum, D. W. Equilibrium, Kinetic, and UV-Spectral Characteristics of Aqueous Bromine, Chloride, Bromine and Chlorine Species. Inorg. Chem. 1994, 33, 5872−5878. (22) Tóth, Z.; Fábián, I. Kinetics and Mechanism of the Initial Phase of Bromine−Chlorite Ion Reaction in Aqueous Solution. Inorg. Chem. 2000, 39, 4608−4614. (23) Horváth, A. K.; Nagypál, I.; Epstein, I. R. Three Autocatalysts and Self-Inhibition in a Single Reaction: A Detailed Mechanism of the Chlorite−Tetrathionate Reaction. Inorg. Chem. 2006, 45, 9877−9883. (24) Ensafi, A. A.; Rezaei, B.; Nouroozi, S. Highly Selective Spectrophotometric Flow-Injection Determination of Trace Amounts of Bromide by Catalytic Effect on the Oxidation of m-Cresolsulfonephtalein by Periodate. Spectrochim. Acta, Part A 2004, 60, 2053− 2057. (25) Bagherian, G.; Chamjangali, M. A.; Amin, A. H.; Ameri, S. Selective Kinetic-Spectrophotometric Determination of Trace Amounts of As(III) Based on an Induction Period. Eur. J. Anal. Chem. 2012, 7, 78−88. (26) Schmitz, G. Kinetics and Mechanism of the Iodate−Iodide Reaction and Other Related Reactions. Phys. Chem. Chem. Phys. 1999, 1, 1909−1914. (27) Barton, A. F. M.; Wright, G. A. Kinetics of the Iodate−Iodide Reaction: Catalysis by Carboxylate and Phosphate Ions. J. Chem. Soc. A 1968, 2096−2103. (28) Crouthamel, C. E.; Hayes, A. M.; Martin, D. S. Ionization and Hydration Equilibria of Periodic Acid. J. Am. Chem. Soc. 1951, 73, 82− 87. (29) Horváth, A. K.; Nagypál, I.; Peintler, G.; Epstein, I. R.; Kustin, K. Kinetics and Mechanism of the Decomposition of Chlorous Acid. J. Phys. Chem. A 2003, 107, 6966−6973. (30) Schmitz, G.; Rooze, H. Mecanisme des Reactions du Chlorite et du Dioxyde de Chlore. 5. Cinetique de la Reaction Chlorite−Bromure. Can. J. Chem. 1987, 65, 497−501.

ASSOCIATED CONTENT

S Supporting Information *

Equilibrium data and molar absorbances of bromine and the tribromide ion as well as the result of the individual fit of each kinetic curve are provided. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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REFERENCES

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Kinetics and Mechanism of the Oxidation of Bromide by Periodate in

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