Article pubs.acs.org/IC
Kinetics and Mechanism of the Chlorite−Periodate System: Formation of a Short-Lived Key Intermediate OClOIO3 and Its Subsequent Reactions Nóra Baranyi, György Csekő, László Valkai, Li Xu, and Attila K. Horváth* Department of Inorganic Chemistry, University of Pécs, Ifjúság útja 6, H-7624 Pécs, Hungary ABSTRACT: The chlorite−periodate reaction has been studied spectrophotometrically in acidic medium at 25.0 ± 0.1 °C, monitoring the absorbance at 400 nm in acetate/acetic acid buffer at constant ionic strength (I = 0.5 M). We have shown that periodate was exclusively reduced to iodate, but chlorite ion was oxidized to chlorate and chlorine dioxide via branching pathways. The stoichiometry of the reaction can be described as a linear combination of two limiting stoichiometries under our experimental conditions. Detailed initial rate studies have clearly revealed that the formal kinetic orders of hydrogen ion, chlorite ion, and periodate ion are all strictly one, establishing an empirical rate law to be d[ClO2]/dt = kobs[ClO2−][IO4−][H+], where the apparent rate coefficient (kobs) was found to be 70 ± 13 M−2 s−1. On the basis of the experiments, a simple four-step kinetic model with three fitted kinetic parameters is proposed by nonlinear parameter estimation. The reaction was found to proceed via a parallel oxygen transfer reaction leading to the exclusive formation of chlorate and iodate as well as via the formation of a short-lived key intermediate OClOIO3 followed by its further transformations by a sequence of branching pathways.
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INTRODUCTION
participate in the periodate−bromide and chlorate−bromide reactions.6,9 It is easily seen that, in previous chlorite-based reaction systems, chlorite ion has been used as an oxidant. The redox potential of the IO4−/IO3− couple is reported to be around +1.05 V within the pH range of 4−5; thus, periodate is indeed a stronger oxidizing agent than chlorite ion under these experimental conditions.15 As a comparison the redox potentials of ClO2/ClO2− and of the ClO3−/ClO2− couple are +0.954 and +0.916 V, respectively;16 thus, when aqueous periodate and chlorite solutions are mixed, periodate is expected to be reduced while chlorate and chlorine dioxide are anticipated as products from the chlorine part. Herein, we report a detailed kinetic study of the title system for the first time via monitoring the absorbance at 400 nm, where only chlorine dioxide absorbs the light, and we suggest a possible mechanism to explain the most important features of the experimental curves.
Redox transformation reactions of oxyhalogen species have been studied extensively ever since chlorite-based chemical oscillators were shown to comprise the largest family of chemical oscillators.1 As a result the last couple of decades have witnessed intense studies of various oxyhalogen-based systems via monitoring the reactants, intermediates, and/or products by different spectroscopic techniques such as UV−vis2 and stopped-flow technique3 in order to establish reliable kinetic models and to obtain thorough insight into their mechanisms. The list, of course, does not end here; further examples of such systems as the bromate−chlorite,4 hypochlorous acid−chlorite,5 chlorate−bromide,6 chlorate−chloride,7 iodide−iodate,8 and periodate−bromide9 reactions may also be provided after a survey of the literature. Checking the reported mechanisms of these systems has revealed that crucial short-lived intermediates (halogen oxides with different compositions) are considered to be involved for explanation of the kinetic behaviors. As classical examples of these intermediates,10,11 Cl2O2 and I2O2 were introduced to explain the diverse kinetic features of the oscillatory reactions of chlorite12 as well as that of the longknown Dushman reaction.8,13 Moreover, at strongly acidic conditions, Cl2O2 was recently taken into consideration as a short-lived transient during the course of the slow chloride− chlorate reaction.7 Cl2O3 was established as a key intermediate in the decomposition of chlorous acid in driving the reaction to stoichiometrically different pathways.14 Very recently mixed halogen oxides (BrIO3 and BrOClO) were reported to © XXXX American Chemical Society
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EXPERIMENTAL SECTION
Materials and Buffers. Commercially available sodium chlorite was recrystallized as described previously,5 and its purity was found to be >99% by iodometric titration. All of the other chemicals (potassium periodate, acetic acid, sodium acetate, and sodium sulfate) were of the highest purity commercially available and were used without further purification. The stock solutions were prepared from twice ionReceived: December 7, 2015
A
DOI: 10.1021/acs.inorgchem.5b02836 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry exchanges and double-distilled water. The pH of the stock solutions was regulated between 3.55 and 4.55 by acetic acid/acetate buffer, taking the pKa of acetic acid as 4.55.17 The concentration of acetate was kept constant at 0.167 M, and the pH was adjusted by the necessary amount of acetic acid. The ionic strength was adjusted to 0.5 M by using sodium sulfate. Kinetics and mechanism of the title reaction were studied at the following concentration ranges: [ClO2−]0 = 0.2−6.8 mM and [IO4−]0 = 0.5−10.4 mM. All these experiments were carried out at 25.0 ± 0.1 °C. Methods and Instrumentation. The reaction was followed by a Zeiss S600 diode array spectrophotometer in the visible range. The deuterium lamp was switched off in order to avoid the well-known photochemical decomposition of chlorine dioxide.14,18 The kinetic measurements were carried out in a standard quartz cuvette equipped with a magnetic stirrer and Teflon cap having 1 cm optical path. The cuvette was carefully sealed with Parafilm at the Teflon cap to minimize the loss of chlorine dioxide. The periodate solution, the buffer components, and the sodium sulfate (if necessary) solution were delivered from a pipet first. The reaction was started with addition of the necessary amount of chlorite solution from a fast delivery pipet. The spectra of the reacting solution were recorded up to 80 000 s at the wavelength range of 400−800 nm. To check the end-products of the reaction, Raman spectroscopy was performed by a NXR FTRaman spectrometer. Data Treatment. The only absorbing species in the visible range was found to be chlorine dioxide; hence, simultaneous evaluation of the kinetic curves was carried out at a single wavelength (400 nm) by the program package ZiTa/ChemMech,19 where the molar absorbance of this species was found to be 571.5 M−1 cm−1. Each original kinetic run contained more than 900 absorbance−time 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.20 Altogether, almost 3200 experimental points from 62 kinetic series were used for the simultaneous evaluation. Our quantitative criterion for an acceptable fit minimizing the average deviation between the measured and calculated absorbances was to obtain 0.005 au, which is close to the experimentally achievable limit of error.
Figure 1. Raman spectra of the vacuum-evaporated reaction mixture. Conditions: [IO4−]0 = 10.9 mM, [ClO2−]0 = 13.8 mM. The reaction mixture was acidified by perchloric acid ([HClO4]0 = 1 mM).
Table 1. Yield of Chlorine Dioxide in the Periodate− Chlorite Reaction in Representative Examples
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RESULTS Stoichiometry. Under our experimental condition, chlorine dioxide and chlorate ion were found to be the chlorinecontaining products. Reduction of periodate cannot lead to the formation of iodide ion because the well-known reactions of iodide with both reactants21,22 would have led to the formation of iodine. In agreement with this characteristic, the visible absorption spectrum of iodine cannot be observed by the end of the reaction. Consequently, periodate ion is reduced exclusively to iodate. Furthermore, chlorine dioxide can easily be detected by UV−vis spectroscopy. To confirm the formation of chlorate and iodate, we performed Raman spectroscopic studies. These investigations revealed that chlorate and iodate were indeed products of the title reaction in excess of periodate and chlorite ions. A representative Raman spectrum is shown in Figure 1. Quantitative analysis of ClO2 yield was also obtained from the measured absorbance at 400 nm in excess. If periodate was employed in a stoichiometric excess, then [ClO3−]∞ + [ClO2]∞ = [ClO2−]0 from the chlorine mass balance; hence, the data from the last column of Table 1 can be obtained easily by A∞/ (εClO2[ClO2−]0). If, however, chlorite was applied in a stoichiometric excess from the electron conservation balance, we may obtain [ClO3−]∞ + [ClO2]∞ = [IO4−]0 + [ClO2]∞/2, from which the last column of Table 1 can also be calculated easily. As can be seen, these data suggest that yield of chlorine dioxide strongly depends on the initial concentration of
[IO4−]0/mM
[ClO2−]0/mM
pH
A∞ (400 nm)
[ClO2 ]∞ ([ClO2 ]∞ + [ClO− 3 ]∞ )
1.5 1.5 1.5 1.5 10.4 7.5 4.5 3.0 1.5 1.5 1.5 1.5 1.5 1.5 1.5 1.5
0.3 0.5 0.7 1.5 1.5 1.5 1.5 1.5 0.27 1.5 1.5 6.75 4.5 5.8 3.87 2.58
3.55 3.55 3.55 3.55 3.55 3.55 3.55 3.55 4.15 3.75 3.95 3.55 3.55 4.15 4.15 4.15
0.057 0.110 0.165 0.37 0.405 0.40 0.405 0.405 0.045 0.365 0.36 0.89 0.75 0.77 0.60 0.50
0.33 0.38 0.41 0.43 0.47 0.467 0.47 0.47 0.33 0.42 0.42 0.68 0.61 0.62 0.52 0.45
chlorite. Simultaneously the calculated result appears to establish that the yield is basically independent of the pH and [IO4−]0. On the basis of these investigations, the stoichiometry of the reaction can be described by an appropriate linear combination of the following equations: 2ClO2− + IO4 − + 2H+ → 2ClO2 + IO3− + H 2O −
−
ClO2 + IO4 →
ClO3−
+
IO3−
(1) (2)
Initial Rate Studies. The kinetic curves were measured at 400 nm, where the main absorbing species was chlorine dioxide. This provides a direct possibility to determine the initial rate of the reaction at this wavelength. Figure 2 shows the results of the initial rate studies. As can be seen, it is clear that the formal kinetic orders of hydrogen ion, chlorite ion, and periodate ion are all strictly one. This result appears to establish the following empirical rate equation for the formation of chlorine dioxide in the case of the title reaction: B
DOI: 10.1021/acs.inorgchem.5b02836 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry
for the detailed calculation, but it is not a central part of the proposed model. The average deviation was found to be 0.0058 absorbance units, indicating a sound description of the measured curves by the proposed model. Kinetic parameters of the proposed model are summarized in Table 2. Figures 3, 4, Table 2. Fitted and Fixed Rate Coefficients of the Proposed Kinetic Modela
−
−
IO4 + ClO2 →
IO3−
+
Figure 3. Measured (filled circles) and calculated (solid lines) absorbance−time curves at [IO4−]0 = 1.5 mM and pH = 4.14 with different initial chlorite concentrations. [ClO2−]0/mM = 5.8 (black), 3.9 (blue), 2.6 (green), 1.3 (cyan), 0.6 (red), 0.4 (magenta), and 0.3 (brown).
and 5 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.
(R1) (R2)
OClOIO3 + H 2O → ClO3− + IO3− + 2H+
(R3)
OClOIO3 + ClO2 → 2ClO2 +
k1 = 37.7 ± 0.3 M−2 s−1 k2 = 38.5 ± 0.1 M−2 s−1 k3 ≥ 1 s−1b k4/k3 = 4409 ± 30 M−1 [c]
(E1)
ClO3−
IO4 − + ClO2− + 2H+ → OClOIO3 + H 2O
−
parameter value
k1[IO4−][ClO2−][H+] k2[IO4−][ClO2−][H+] k3[OClOIO3] k4[OClOIO3][ClO2−]
(3)
Hence, kobs can be obtained easily from the calculated initial rates, dividing them by the initial concentration of the reactants as well as by the molar absorbance of chlorine dioxide for each kinetic run individually. From these calculations we obtained kobs = 70 ± 13 M−2 s−1. Later we shall see that the relatively large deviation of the observed rate coefficient stems from the slight dependence of kobs on [ClO2−]0. Proposed Kinetic Model. On the basis of the experimental observations, the proposed kinetic model must include the following species: the reactants (periodate and chlorite), the products (chlorate, iodate, and chlorine dioxide), the solvent, the H+, the OH−, and a conceivable intermediate like OClOIO3. At the beginning of the fitting process, the kinetic model contained all the possible mono- and bimolecular reactions between the previously mentioned species. During the fitting procedure, if a rate coefficient became insensitive to the average deviation, the given step was eliminated from the kinetic model. This method was used successfully in a number of our earlier works.14,20 After a long systematic reduction procedure, the following kinetic model was finally proposed: CH3COOH ⇌ H+ + CH3COO−
rate equation
R1 R2 R3 R4
a No error indicates that the given parameter is fixed during the calculation process. bOnly a lower limit can be determined. cOnly k4/ k3 can be determined from our measurements.
Figure 2. Initial rate studies at different conditions. In the case of the black curves, the title of x-axis c0 corresponds to [IO4−]0 at pH = 3.55 and [ClO2−]0 = 1.5 mM (•); and pH = 4.15 and [ClO2−]0 = 1.29 mM (deg). In the case of the blue curves (the y-axis is shifted by +0.2), the title of x-axis c0 corresponds to ClO2− at pH = 3.55 and [IO4−]0 = 1.5 mM (•); and pH = 4.15 and [IO4−]0 = 1.5 mM (deg). In the case of the red curves, the title of x-axis c0 corresponds to [H+] at [IO3−]0 = 1.5 mM and [ClO2−]0 = 1.5 mM (•) (the x-axis is shifted by +1.0); and [IO4−]0 = 6.3 mM and [ClO2−]0 = 1.2 mM (deg) (the x-axis is also shifted by +1.0 and the y-axis is shifted by +0.3).
d[ClO2 ] = kobs[ClO−2 ][IO−4 ][H+] dt
step
IO3−
(R4)
Figure 4. Measured (filled circles) and calculated (solid lines) absorbance−time curves at pH = 4.15 and [ClO2−]0 = 1.4 mM with different initial periodate concentrations. [IO4−]0/mM = 1.0 (black), 1.5 (blue), 2.0 (green), 3.0 (cyan), and 4.5 (red). [ClO2−]0 = 0.16 mM and pH = 4.15 with different initial periodate concentrations. [IO4−]0/ mM = 0.3 (magenta), 0.5 (brown), 0.8 (gray), and 1.0 (yellow).
The rapid protonation process (eq E1) was taken into account with known equilibrium constant17 to follow the slight change in pH during the reaction. This acid dissociation equilibrium may be regarded as an auxiliary process necessary C
DOI: 10.1021/acs.inorgchem.5b02836 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry
Step R3 is a fast hydrolysis of OClOIO3, leading to the formation of chlorate and iodate ions. The individual rate coefficient of this reaction cannot be determined from our measurements; only k4/k3 could be calculated. We therefore fixed the value of k3 to be its lower limit, 1 s−1. One might argue that this route can be eliminated from the proposed model because an alternative pathway exists to produce chlorate ion (see step R1). We have thus performed an additional calculation process with elimination of this step. This calculation provided an average deviation of 0.022 au, from which we concluded that this step is also necessary to describe quantitatively our kinetic data. The reason for the failure of this simplified model can be understood in terms of the data summarized in Table 1. As can be seen, elimination of k3 means that the ratio [ClO2]∞/[ClO3−]∞ depends just on the ratio of k2 and k1 but would not depend on [ClO2−]. Because this dependence is clearly established in Table 1, the role of this pathway is confirmed. Step R4 is also an essential pathway for the formation of chlorine dioxide analogous to the case of hypochlorous acid− chlorite reaction.5 Because the rate of this reaction is directly proportional to [ClO2−], a shift in the yield of chlorine dioxide can easily be taken into account. As was mentioned earlier, the individual rate coefficient of this reaction cannot be calculated; only the ratio of k4/k3 = 4409 ± 30 M−1 was determined. Formal Kinetics. Because OClOIO3 is a short-lived intermediate, it provides us a possibility to derive the rate equation for the formation of chlorine dioxide. Applying steadystate approximation for OClOIO3, one can easily obtain the following expressions:
Figure 5. Measured (filled circles) and calculated (solid lines) absorbance−time curves at different pHs. [ClO2−]0 = 1.35 mM and [IO4−]0 = 1.5 mM. pH = 3.55 (black), 3.75 (blue), 3.95 (green), 4.15 (cyan), 4.35 (red), and 4.55 (magenta).
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DISCUSSION Step R1 is a direct oxygen transfer process between the reactants periodate and chlorite ions. This is one of the ratedetermining steps of the overall process that is responsible for the relatively low yield of chlorine dioxide. As was visualized earlier (see Table 1), even in large excess of chlorite a significant amount of chlorate forms (>30% of the chlorine is transformed into chlorate), so the role of this step is strongly established in the proposed model. To confirm directly its importance, an additional fitting procedure was performed with elimination of this step. The average deviation of this fit was found to be 0.04 au, indicating the necessity of this step in the proposed model. Step R2 is the other rate-determining step of the proposed model, producing OClOIO3, a key intermediate of the reaction. Similar oxyhalogen intermediates have already been wellestablished in the case of the Dushman reaction,8 as well as those of the chlorite−hypochlorous acid,5 chlorate−chloride,7 periodate−bromide,9 and chlorate−bromide6 reactions. By an analogy, we therefore propose here OClOIO3. The most important difference about the formation of this species is that it forms in an irreversible reaction. We have also tried to treat this process in our model as a rapidly established equilibrium shifted far to the left but this would lead to a contradiction with our kinetic data if further reactions of OClOIO3 are steps R3 and R4. The reason is simple if step R2 is a fast pre-equilibrium; then combining this step with step R4 would lead to a formal kinetic order of two for chlorite ion. Our kinetic data, however, have provided no evidence to support this fact. Purely firstorder kinetics with respect to chlorite ion suggests that, if OClOIO3 forms in a rapidly established equilibrium, then the intermediate OClOIO 3 first must proceed via a ratedetermining step (possibly a homolytic or a heterolytic bondcleavage reaction) whose rate does not depend on [ClO2−]. Such processes, however, seem to be so unfavorable thermodynamically that we propose that either this step is an irreversible process (this is, of course, much more preferred) or it is a very slowly established equilibrium and the backward reaction essentially does not have any influence on the time scale of the reaction. Finally, this route plays a crucial role because formation of chlorine dioxide takes place exclusively by this way, so this step cannot be eliminated from the proposed model.
[OClOIO3] =
k 2[IO4 −][ClO2−][H+] k 3 + k4[ClO2−]
(4)
From the kinetic model it is easy to conclude that the following differential equation is valid for the formation of chlorine dioxide: d[ClO2 ] = 2k4[OClOIO3][ClO2−] dt
(5)
By substituting eq 4 into eq 5 followed by some algebraic rearrangement, the following expression is derived: k4[ClO2−] d[ClO2 ] = 2k 2[IO4 −][ClO2−][H+] k 3 + k4[ClO2−] dt = kobs[IO4 −][ClO2−][H+]
(6)
Considering the smallest (0.2 mM) and the largest (6.8 mM) initial chlorite concentrations and the values of k2, k3, and k4 obtained from the nonlinear parameter estimation, we obtain the value of kobs within the range 36.2−74.7 M−2 s−1, which is in very good agreement with 70 ± 13 M−2 s−1 calculated from the initial rates. At the same time, it explains pretty much as well the relatively high average deviation of kobs obtained in the case of the individual evaluations. Finally, a word is also in order here to discuss possible alternative pathways instead of invoking the presence of the short-lived intermediate OClOIO3. Very recently, Pan and his co-workers confirmed that the sulfite−chlorine dioxide reaction proceeds via parallel oxygen and electron transfer routes.23 Considering this opportunity here, we also tried to exchange the OClOIO3-driven route to the following sequence of reactions: D
DOI: 10.1021/acs.inorgchem.5b02836 Inorg. Chem. XXXX, XXX, XXX−XXX
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Inorganic Chemistry k 2′ , k −′ 2
ClO2− + IO4 − HoooooI ClO2 + IO4 2 − k 3′
IO4 2 − + ClO2− + 2H+ → ClO2 + IO3− + H 2O
However, a notably higher average deviation found rather supports the fact that the reaction is more likely to proceed via the formation of OClOIO3.
(R2′)
■
(R3′)
and the measured absorbance−time series were evaluated by steps R1, R2′, and R3′ refining parameters k1, k′2, k′−2, and k′3. We found a total correlation between k′−2 and k′3, meaning that k1 = 26.9 ± 0.2 M−2 s−1, k2′ = 36.9 ± 0.1 M−2 s−1, and k3′ /k−2 ′ = 0.74 ± 0.02 can be determined from our data with an acceptable average deviation of 0.0084 au. It is, however, ∼50% higher than the best average deviation found by the proposed model from which we concluded that, according to our data, it is more likely that the reaction proceeds via the intermediate OClOIO3. Because the proposed model contains the hydrolysis of OClOIO3 and the reaction of this species with chlorite ion, this behavior seems to be reminiscent of the case of Cl2O2 and I2O2. It is generally well-known that subsequent reactions of Cl2O2 and I2O2 are often subject to be affected by the concentration of the buffer components as well. Performing extra experiments when only the concentration of acetate was changed between 0.038 and 0.265 mM and at pH = 3.55, we found absolutely no effect of the concentration of buffer components on the formation of chlorine dioxide, which is quite different from the case of the reactions of Cl2O2 and I2O2. In those systems, however, these species form in a rapid preequilibrium established instantaneously but shifted far to the left, meaning that kinetic features of the overall reaction should also depend on the characteristics of their subsequent reactions. In the present system, however, the rate-determining step is the formation of OClOIO3; hence, the kinetic feature of their subsequent reactions is not manifested in the overall kinetics. Therefore, lack of buffer assistance in the present case can easily be explained by the fact that the OClOIO3−chlorite direct reaction takes place after the rate-limiting steps.
AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS Financial support of the Hungarian Research Fund OTKA Grant No. K116591 is gratefully acknowledged. Financial support for the “Environmental industry related innovative trans- and interdisciplinary research team development in the University of Pécs knowledge base” Project (SROP-4.2.2.D-15/ 1/KONV-2015-0015) is gratefully acknowledged. The present scientific contribution is dedicated to the 650th anniversary of the foundation of University of Pécs, Hungary.
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REFERENCES
(1) Epstein, I. R.; Orbán, M. In Oscillations and Traveling Waves in Chemical Systems; Field, R. J., Burger, M., Eds.; Wiley: New York, 1985; pp 258−286. (2) Wang, X.; Kelley, M. D.; Cooper, J. N.; Beckwith, R. C.; Margerum, D. W. Inorg. Chem. 1994, 33, 5872−5878. (3) Cortes, C. E. S.; Faria, R. B. Inorg. Chem. 2004, 43, 1395−1402. (4) Schmitz, G.; Rooze, H. Can. J. Chem. 1988, 66, 231−235. (5) Peintler, G.; Nagypál, I.; Epstein, I. R. J. Phys. Chem. 1990, 94, 2954−2958. (6) Sant’Anna, R. T. P.; Faria, R. B. Inorg. Chem. 2015, 54, 10415− 10421. (7) Sant’Anna, R. T. P.; Santos, C. M. P.; Silva, G. P.; Ferreira, R. J. R.; Oliveira, A. P.; Cortes, C. E. S.; Faria, R. B. J. Braz. Chem. Soc. 2012, 23, 1543−1550. (8) Schmitz, G. Phys. Chem. Chem. Phys. 1999, 1, 1909−1914. (9) Szél, V.; Csekő , G.; Horváth, A. K. J. Phys. Chem. A 2014, 118, 10713−10718. (10) Taube, H.; Dodgen, H. J. Am. Chem. Soc. 1949, 71, 3330−3336. (11) Bray, W. C. J. Am. Chem. Soc. 1930, 52, 3580−3586. (12) Rábai, G.; Orbán, M. J. Phys. Chem. 1993, 97, 5935−5939. (13) Dushman, S. J. Phys. Chem. 1904, 8, 453−482. (14) Horváth, A. K.; Nagypál, I.; Peintler, G.; Epstein, I. R.; Kustin, K. J. Phys. Chem. A 2003, 107, 6966−6973. (15) Berka, A.; Vulterin, J.; Zyka, J. Newer Redox Titrants: International Series of Monographs in Analytical Chemistry; Pergamon Press: Headington Hill Hall, Oxford, 1965; pp 66−75. (16) Lide, D. L. CRC Handbook of Chemistry and Physics, 71st ed.; CRC Press: Boca Raton, FL, 1990. (17) IUPAC. Stability Constant Database; Royal Society of Chemistry: London, 1992−1997. (18) Stanbury, D. M.; Figlar, J. N. Coord. Chem. Rev. 1999, 187, 223− 232. (19) Peintler, G. ZiTa, version 5.0; A Comprehesive Program Package for Fitting Parameters of Chemical Reaction Mechanism; Attila József University: Szeged, Hungary, 1989−1998. (20) Horváth, A. K.; Nagypál, I.; Epstein, I. R. Inorg. Chem. 2006, 45, 9877−9883. (21) Lengyel, I.; Li, J.; Kustin, K.; Epstein, I. R. J. Am. Chem. Soc. 1996, 118, 3708−3719. (22) Horváth, A. K. J. Phys. Chem. A 2007, 111, 890−896. (23) Pan, C.; Gao, Q.; Stanbury, D. M. Inorg. Chem. 2016, 55, 366− 370.
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CONCLUSION Our present report may be treated as a comprehensive effort to unravel the kinetics and mechanism of the chlorite−periodate reaction. Detailed initial rate studies have clearly revealed that the formal kinetic orders of hydrogen ion, chlorite ion, and periodate ion are strictly one. The study presented here also indicated that the reaction can be described by two limiting stoichiometries, and the actual stoichiometry of the reaction can be established by an appropriate linear combination of these equations in the concentration ranges studied. Furthermore, a key intermediate OClOIO3 was proposed as well to describe the kinetic data. It is also shown that this intermediate disappears via two different pathways: (1) hydrolysis of this compound to produce chlorate and iodate exclusively as well as (2) its direct reaction with chlorite ion leading to the formation of chlorine dioxide. These reactions altogether constitute a four-step kinetic model with three fitted kinetic parameters that is able to describe soundly the most important characteristics of the measured kinetic curves. Lack of buffer assistance found experimentally confirms that direct reaction between OClOIO3 and substrate (chlorite) occurs after the rate-limiting step, which seems to be unique compared to the cases of other oxyhalogen key intermediates. We have also shown that an alternative model is also capable of describing the kinetic data adequately. In this model the formation of OClOIO3, along with its further transformations, is substituted by a sequence of electron-transfer reactions. E
DOI: 10.1021/acs.inorgchem.5b02836 Inorg. Chem. XXXX, XXX, XXX−XXX
