Iodine(+1) Reduction by Hydrogen Peroxide - Springer Link

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The reduction of iodine(+1) by hydrogen peroxide is also important as a part of the Bray–Liebhafsky oscil lating reaction, that is the decomposition (8) catalyzed.

ISSN 00360244, Russian Journal of Physical Chemistry A, 2009, Vol. 83, No. 9, pp. 1447–1451. © Pleiades Publishing, Ltd., 2009.

CHEMICAL KINETICS AND CATALYSIS

Iodine(+1) Reduction by Hydrogen Peroxide* G. Schmitz Faculté des Sciences Appliquées, Université Libre de Bruxelles, Belgium email: [email protected] Abstract—The iodine(+1) reduction by hydrogen peroxide is catalyzed by different buffers and its rate is a complicated function of the acidity and of the iodide concentration. The seemingly inconsistent published experimental results are reanalyzed and a new kinetic model is proposed. A key step is the catalysis by the buffers of the formation of the intermediate compound IOOH. This model reconciles the previous works. DOI: 10.1134/S0036024409090052

INTRODUCTION Although the reduction of iodine(+1) by hydrogen peroxide has been studied during more than eighty years, there is no consensus in the literature about its kinetics. The first investigations [1–3] dealt with the catalytic decomposition of hydrogen peroxide by iodine. The following mechanism was postulated in neutral or dilute acidic solutions: I2 + H2O IOH + I– + H+, (1) –

I2 + I–

I3 , IO– + H+,

IOH –

I–

IO + H2O2 –

+ H2O + O2,

IO–

I + H2O2

+H2O.

(2) (3) (4) (5)

In more acidic solutions, the mechanism includes also reactions IOH + H2O2 I– + H+ + H2O + O2, (6) I– + H+ + H2O2

IOH + H2O.

(7)

The sum of reactions (4) and (5) or (6) and (7) gives the decomposition. 2H2O2 2H2O + O2. (8) The reduction of iodine(+1) by hydrogen peroxide is also important as a part of the Bray–Liebhafsky oscil lating reaction, that is the decomposition (8) catalyzed by iodate and iodine in acidic solutions [4–9]. This decomposition is the result of the global reactions –

2IO 3 + 2H+ + 5H2O2 I2 + 5H2O2

I2 + 5O2 + 6H2O, –

2IO 3 + 2H+ + 4H2O.

(9) (10)

In reaction (9), hydrogen peroxide acts as a reducing agent but its direct reaction with iodate is much too slow to explain the observed rates. The mechanism is complicated and the reducing action of hydrogen per *The article is published in the original.

oxide is mainly the result of reaction (6). Liebhafsky [3, 10] and Furrow [11] have tried to isolate this reac tion and to measure its rate constant but their values differ by a factor ten. A third group of investigations deals with the importance of the iodine(+1) reduction by hydrogen peroxide in analyzing the iodine behavior after a nuclear reactor accident. 131I is one of the most toxic fission products that would be released in the atmo sphere and the behavior of iodine has been the subject of extensive studies for the nuclear industry [12–16]. An important part of these studies were devoted to the iodine reactions in solution controlling the distribu tion of iodine between volatile compounds (I2 and organic compounds) and nonvolatile compounds – (mainly I– and IO 3 ). Hydrogen peroxide being one of the products of the water radiolysis, the reactions of iodine(+1) with hydrogen peroxide producing iodide or iodate are important processes that would influence the release of 131I in the atmosphere. The present state of our knowledge about the reduction of iodine(+1) by hydrogen peroxide is mainly the result of the works of Liebhafsky [2, 3], of Shiraishi et al. [12, 13] and of Ball and Hnatiw [16, 17]. Liebhafsky (1932) has studied the catalytic decom position (8) in an acetate buffers between pH 4 and 6 and in a phosphate buffers between pH 6 and 8 [2]. The author has concluded that, under his experimen tal conditions, the rate of iodide oxidation (5) is nearly equal to the rate of the iodine(+1) reduction (4). From the observed rates of hydrogen peroxide decomposi tion and the wellknown rate constant k5 [18], k4 = 6 × 109 mol–1 dm3 s–1 was calculated. The obtained val ues were independent on the acidity, the nature of the buffer and the iodide concentration. In another work [3] the author has measured the rate of oxygen produc tion by reaction (4) far from the steady state and has obtained k4 = 3.3 × 109 mol–1 dm3 s–1 in satisfactory agreement with the above value. Both works suggest that the rate is proportional to [IO–] and is indepen

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dent on the acidity, the phosphate concentration and the iodide concentration. It is this simplicity that led Liebhafsky to think that (4) was the rate determining step but further measurements have shown that it was an illusion. Shiraishi et al. (1991) have discovered that phos phate buffers catalyze the iodine(+1) reduction by hydrogen peroxide while citrate buffers do not [12]. For this reason they have performed a new kinetic study in a citrate buffer between pH 5.5 and 7 [13] and found a new term in the rate law, proportional to [I–]. To explain this term, they have added reactions –

IOH + I I2OH– + H2O2



(11)

2I– + H+ + H2O + O2

(12)

I2OH ,

to the mechanism proposed by Liehhafsky. Their value k4 = 7 × 107 mol–1 dm3 s–1 is much lower than Liebhaf sky value. Ball and Hnatiw (2001) [16] have observed devia tions from the rate law of Shiraishi et al. and have decided to perform a more complete kinetic study as a function of acid, hydrogen peroxide, iodide and buffer concentration in three different buffers [17]. On the basis of this very important experimental work they have proposed the mechanism including (11) and reactions: I2OH– + H2O2 IOOH +

OH–

IOOH + B–

I– + IOOH + H2O, I–

+ H2O + O2,

I– + BH + O2,

(13) (14) (15)

where B– denotes the anion of the buffer. This mechanism explains their experimental results but we will see that it is inconsistent with the former works. We can reject proposed mechanisms but we cannot ignore experimental results without a good reason. Thus, we have analyzed the published rate measurements, found that they are compatible and that a modification of the Ball and Hnatiw mechanism can explain them all. ANALYSIS OF THE EXPERIMENTAL RESULTS The observed evolutions of the iodine concentra tion discussed in this work are the result of two reac tions, the reduction of iodine(+1) by hydrogen perox ide and the oxidation of iodide to iodine by reaction (5). Denoting by [I2]t the total iodine(+1) concentration, –

[I2]t = [I2] + [ I 3 ] + [IO–] + [IOH] + [I2OH–], and by kobs [I2]t the global rate of iodine(+1) reduction, the expression of the observed rates is (16) where r5 = k5[I–][H2O2]: –d[I2]t/dt = kobs[I2]t – r5.

(16)

During the reduction of iodine(+1), kobs[I2]t decreases and r5 increases until kobs[I2]t = r5. Then, the system

has reached the steady state of the catalytic decompo sition studied by Liebhafsky. The different authors have presented their experi mental results differently and in order to compare them it is necessary to choose a common presentation. For this reason, we define the following new function Z (mol2 dm–6 s–1). Its values can be calculated using the information published by the different authors and we will see that its variations with the experimental conditions gives a convenient representation of the effects discussed in this work: 14 I ] [ H ] [ I2 ]t Z = [  k obs × 10 . [ H2 O2 ] [ I2 ] –

+ 2

Under the conditions of all the works discussed here the reactions (1)–(3) are quasi at equilibrium and the concentrations [IO–] and [I2OH–] can be neglected in the iodine mass balance. This gives the following rela tions: [IOH][I–][H+] = K1[I2],



[ I 3 ] = K2[I2][I–],

[IO–][H+] = K3[IOH], [I2]t = [I2] (1 + K2[I–] + K1/[I–][H+]). The mechanism proposed by Liebhafsky assumes that kobs [I2]t = k4[IO–][H2O2], where [IO–] = K1K3[I2]/[I–][H+]2. If so, these expressions intro duced in the definition of Z would give Z = k4K1K3 × 1014. However, the experimental values of Z are not constant. The effects of the buffers, the acidity and the iodide concentration on the Z values show the kinetic complications revealed by further works. Kinetics in Citrate Buffers Shiraishi et al. as well as Ball and Hnatiw have con cluded that citrate buffers do not catalyze significantly the iodine(+1) reduction by hydrogen peroxide. On the other hand, Shiraishi et al. have observed a marked effect of the iodide ions, not observed by Ball and Hnatiw. The values given in their figures allow the cal culation of Z and lead to Z = 0.2 + 1.8 × 108[H+][I–]. (17) A careful analysis of the results of Ball and Hnatiw reveals why they have not observed this effect: in the rate expression (16) d[I2]t/dt is proportional to [I2]t only if r5 can be neglected. Shiraishi et al. have ana lyzed their results using the complete expression (16) but Ball and Hnatiw have calculated kobs assuming a simple first order rate law. This approximation was correct for nearly all their measurements but not for their study of the [I–] effect in citrate buffers. At high iodide concentrations the r5 term is important and their kobs values are too low. They did not publish the details of their measurements but the available infor mation suggests that their values corrected for the r5 term would be similar to Shiraishi et al. values.

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Kinetics in Phosphate Buffers Ball and Hnatiw have measured kobs between pH 6.2 and 7.7 with phosphate total concentrations equal to 0.02 or 0.05 mol dm–3 and iodide concentrations equal to 5 × 10–4 or 1 × 10–3 mol dm–3. Figure 1 shows the increase of their reaction rates with [H+][B–]. Iodide has a much smaller effect than the buffer. Figure 1 shows also the Z values calculated from the Liebhafsky measurements between pH 5.8 and 7.2 in 0.2 mol dm–3 phosphate buffer with [I–] between 0.001 and 0.1 mol dm–3. The scatter is important but these val ues are in qualitative agreement with Ball values. They are higher because the phosphate buffer concentration is higher and they seem independent on [I–]. Liebhaf sky measurements suggest that Z reaches a limiting value at high buffers concentrations. For comparison, Fig. 1 shows also the Z value calculated from the mea surements by Shiraishi et al. for the reaction not cata lyzed without effect of iodide (Z = 0.2 mol2 dm–6 s–1). The experimental results of these different works are in good agreement and the kinetic model must explain them all. Kinetics in Barbital Buffers Most of the Ball and Hnatiw results were obtained in barbital buffers. The authors have explained the observed effects of the buffer and of the iodide con centrations by the mechanism (13)–(15) giving the rate law –

k obs [ I 2 ] t = k 13 [ I 2 OH ] [ H 2 O 2 ] –



k 14 [ OH ] + k 15 [ B ]  . ×  – – – k –13 [ I ] + k 14 [ OH ] + k 15 [ B ] With the quasiequilibrium (11) and Kw = [H+][OH–] = 10–14, the corresponding expression of Z is Z (18) + – + k 13 K 1 K 11 × 10 [ I ] [ H ] ( k 14 K w + k 15 [ B ] [ H ] )   . =  – + – + k –13 [ I ] [ H ] + k 14 K w + k 15 [ B ] [ H ] 14

1449

Z 4 3 2 1 0

1

2

1

3

2

4

5

3 4 [H+][B−] × 108

Fig. 1. Comparison of the Z values obtained by different authors. Values calculated from Ball and Hnatiw measure ments in phosphate buffers (1) 0.02 M and (2) 0.05 M [17] and from Liebhafsky measurements in phosphate buffers (3) 0.20 M [2] or (4) [3]; (5) value for the noncatalyzed reaction [13].

tion (5) as indicated above. The expression (18) is also inconsistent with the results of Liebhafsky in phosphate and acetate buffers. The experimental Z values at high acetate or phosphate concentrations shown in Fig. 1 are independent on [I–]. However, the expression (18) predicts the opposite: when the term k15[B–][H+] becomes so large that Z becomes independent on the buffer concentration, Z should become proportional to [I–][H+]. Liebhafsky has mentioned in a footnote ([2], p. 1798) results obtained at a lower phosphate concentration showing “a positive trend with increasing [I–].” Unfortunately he has dis carded these results but this footnote confirms that the effect of [I–] on Z decreases when the buffer concentra tion increases. In summary, Ball and Hnatiw mecha nism is inconsistent with Shiraishi et al. measurements, with Ball and Hnatiw kobs values in citrate buffers cor rected for the effect of reaction (5) and with Liebhafsky measurements.



This expression can be fitted to the author’s experi mental values but is inconsistent with the other works. In citrate buffers the catalytic term k15[B–][H+] can be neglected, the rate constants obtained by Ball and Hnatiw give k14Kw/k–13 = 2 × 10–12 mol2 dm–6 and the expression of Z reduces to –

+



+

Z = K 13 K 1 K 11 k 14 [ I ] [ H ]/ ( [ I ] [ H ] + 2 × 10

– 12

).

The range of [I–][H+] values is 5 × 10–11 to 6 × 10–10 for the Ball and Hnatiw experiments and 5 × 10–11 to 8 × 10–9 for the Shiraishi et al. experiments. In both cases [I–][H+] is larger than 2 × 10–12 and Z should be inde pendent on [I–][H+]. This is in contradiction with Shiraishi et al. results and also with Ball and Hnatiw kobs values if they are corrected for the effect of reac RUSSIAN JOURNAL OF PHYSICAL CHEMISTRY A

PROPOSED MECHANISM The above contradictions can be resolved consider ing the following mechanism where reactions (3) and (11) are quasi at equilibrium: IOH IOH + I IO– + H2O2 IOH + B– + H2O2 I2OH– + H2O2



IO– + H+,

(3)



(11)

I2OH , IOOH + OH–,

(19)

IOOH + BH + OH–, (20) IOOH + I– + H2O,

(13)

IOOH + OH–

I– + H2O + O2,

(14)

IOH + H2O2

IOOH + H2O.

(21)

This mechanism is similar to the one proposed by Ball and Hnatiw but the catalysis of the IOOH decomposi Vol. 83

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SCHMITZ Z 0.3 (b)

(a)

0.2 1 2 3 0.1 1

0

2

0

1

2 [H+] × 107

Fig. 2. Comparison of the calculated Z values (lines) with experimental values in barbital buffers [17]; (a) [I–] = 5 × 10–4 M, [Bar bital] = (1) 1 × 10–2, (2) 5 × 10–3, and (3) 2.5 × 10–3 M; (b) [Barbital] = 5 × 10–3 M, [I–]: (1) 1 × 10–3, (2) 5 × 10–4, and (3) 2.5 × 10–4 M.

tion (15) is replaced with the catalysis of its forma tion (20). The noncatalyzed reaction (21) can be neglected under the experimental conditions of the studies discussed here but is probably important in nonbuffered acidic solutions. The reactions (19), (20), and (13) are kinetically equivalent to the catalysis of (21) by OH–, B–, and I–, respectively. This mecha nism gives the rate law: k obs [ I 2 ] t = k 14 [ H 2 O 2 ] –





k 19 [ O ] + k 20 [ IOH ] [ B ] + k 13 [ I 2 OH ]  . ×  – + k 14 + k –19 + k –20 [ BH ] + k –13 [ I ] [ H ]/K w The rate constants are related by the condition of internal consistency of reaction mechanisms [19]. As (11) + (13) = (21) we must have K11K13 = K21 or K11k13/k–13 = K21. Similar relations for reactions (19) and (20) give k 13 K k 19 K k 20 K 21 = K 11   = 3   = B  , k –13 K w k –19 K w k –20

(22)

where KB = [H+][B–]/[HB]. These relations intro duced in the rate law lead to 14

Z = 10 k 14 K 21 K 1 K w – + – + k 19 K 3 + k 13 K 11 [ I ] [ H ] + k 20 [ B ] [ H ] (23)   . ×  – + – + k 14 K 21 K w + k 19 K 3 + k 13 K 11 [ I ] [ H ] + k 20 [ B ] [ H ]

Let us show firstly that this expression is in qualitative agreement with the experimental results if k14K21Kw is

much larger than k19K3 + k13K11[I–][H+] giving the approximate expression Z ∼ k 14 K 21 K 1 –

+



+

k 19 K 3 + k 13 K 1 [ I ] [ H ] + k 20 [ B ] [ H ] ×    . – + k 14 K 21 K w + k 20 [ B ] [ H ]

(24)

In citrate buffers the catalytic term k20[B–][H+] can be neglected and this expression reduces to Z = K1 (k19K3 + k13K11[I–][H+])/Kw in accordance with equation (17) obtained by Shiraishi et al. The expression (24) explains also the maximum value of Z observed by Lie bhafsky in concentrated buffers. When k20[B–][H+] is very large Z approaches k14K21K1 and [I–] has no effect. When k20[B–][H+] is smaller, the term k13K11[I–][H+] explains the small effect of [I–] observed. Thus, the simplified Eq. (24) explains qualitatively all the obser vations. To show that the agreement is also quantita tive, we have estimated the rate constants using the complete Eq. (23). For numerical reasons, this was performed in two steps. A first analysis of all the differ ent works shows that k14K21K1 must be about 4 s–1. Then, keeping this value, the adjustment of the other rate constants to Shiraishi et al. results gives k19K3K1 = 2 × 10–15, k13K11K1 = 2.3 × 10–6 and k20 very small, as expected in citrate buffers. The adjustment to Ball and Hnatiw results in barbital buffers gives k19K3K1 = 1.1 × 10–15, k13K11K1 = 2.8 × 10–6, in fair agreement with Shiraishi et al. values, and k20K1 = 9 × 10–6. These val ues depend somewhat on the chosen KB value and we have used pKB = 7.5 for 0.2 mol dm–3 ionic strength. Figure 2 shows the agreement between the experimen tal and calculated values. The rate constants of the steps (13), (14), and (19) can be estimated as follow. The equilibrium constants 0 0 0 K 1 = 5.3 × 10–13, K 3 = 2.3 × 10–11, and K 11 = 320 are

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CONCLUSIONS Liebhafsky had used two different experimental methods in different acetate and phosphate buffers and had obtained about the same value of Z. This let him to conclude logically that reaction (4) was the rate determining step. Actually, he had observed a special kind of homogeneous catalysis, a catalytic reaction whose rate can be independent on the nature and on the concentration of the catalyst. The proposed mech anism explains Liebhafsky results: when the action of the buffer is very effective, reaction (20) is quasi at equilibrium and the rate determining step is (14). We have shown that in this case Z is close to k14K21K1, independent on [B–] and also on KB. The work of Lie bhafsky is instructive for another reason: it shows that we can never discard experimental results just because they seem abnormal. Liebhafsky had observed that at low phosphate concentrations the values of Z were lower but had discarded these results because they seemed abnormal. Actually, they were the most infor mative results revealing the complication of the kinet ics of the iodine(+1) reduction by hydrogen peroxide. It is only sixty years later that Shiraishi et al. have dis covered the catalytic effect of the buffers and the com

RUSSIAN JOURNAL OF PHYSICAL CHEMISTRY A

plicated effect of the iodide concentration. Recently, Ball and Hnatiw have measured more precisely these effects but their mechanism does not explain all the former observations. The mechanism we propose explains all the well established experimental facts but we do not pretend that this is the end of the story. REFERENCES 1. E. Z. Abel, Z. Phys. Chem. 96, 1 (1920); Z. Phys. Chem. 136, 16 (1928). 2. H. A. Liebhafsky, J. Amer. Chem. Soc. 54, 1792 (1932). 3. H. A. Liebhafsky, J. Amer. Chem. Soc. 54, 3499 (1932). 4. G. Schmitz, J. Chim. Phys. 84, 957 (1987). 5. Lj. KolarAni c  and G. Schmitz, J. Chem. Soc., Fara day Trans. 88, 2343 (1992). 6. Lj. KolarAni c , Z . Cupi c , S. Ani c , and G. Schmitz, J. Chem. Soc., Faraday Trans. 93, 2147 (1997). 7. G. Schmitz, Phys. Chem. Chem. Phys. 1, 4605 (1999). 8. G. Schmitz, Phys. Chem. Chem. Phys. 3, 4741 (2001). 9. G. Schmitz, Lj. KolarAni c , S. Ani c , T. Grozdi c , and V. J. Vukojevi c , Phys. Chem. A 110, 10361 (2006). 10. I. Matsuzaki, R. Simic, and H. A. Liebhafsky, Bull. Chem. Soc. Jpn. 45, 3367 (1972). 11. S. Furrow, J. Phys. Chem. 91, 2129 (1987). 12. K. Ishigure, H. Shiraishi, H. Okuda, and N. Fujita, Radiat. Phys. Chem. 28, 601 (1986). 13. H. Shiraishi, H. Okuda, Y. Morinaga, and K. Ishigure, in Proc. of the 3rd CSNI Workshop Iodine Chem. Reactor Safety, 1991, JAERIM–92–012 (Japan, 1992), p. 152. 14. J. C. Wren and J. M. Ball, Radiat. Phys. Chem. 60, 577 (2001). 15. B. Cl e ment, L. Cantrel, G. Ducros, F. Funke, L. Her  ranz, A. Rydl, G. Weber, and C. Wren, State of the Art Report on Iodine Chemistry, Nuclear Energy Agency, NEA/CNSI/R(2007)1. 16. J. M. Ball, J. B. Hnatiw, and H. E. Sims, in Proc. of the 4th CSNI Workshop Iodine Chem. Reactor Safety, 1996, Ed. by S. Guntay (Switzerland, 1997), p. 169. 17. J. M. Ball and J. B. Hnatiw, Can. J. Chem. 79, 304 (2001). 18. H. A. Liebhafsky and A. Mohammad, J. Amer. Chem. Soc. 55, 3977 (1933). 19. G. Schmitz, J. Chem. Phys. 112, 10714 (2000). 20. G. Schmitz, Int. J. Chem. Kinet. 36, 480 (2004). ˆ

well known at zero ionic strength [20]. The values of [H+] given by Shiraishi and by Ball are actually those of 10–pH or a H+ and a correction for the ionic strength must be applied only to the monovalent negative ions. With a value 0.8 for their activity coefficients, k13K11K1 = 2.8 × 10–6 gives k13 = 1.3 × 104 mol dm–3 s–1 and k19K3K1 = 1.1 × 10–15 gives k19 = 5.8 × 107 mol dm–3 s–1. The rate constants in the backward directions could be calculated using the relations (22) if the equilibrium constant K21 was known. It is not but some indirect estimation can be obtained. From k14K21K1 = 4 s–1 we calculate k14K21 = 6 × 1012 and note that k14 cannot be larger than about 5 × 109, the value for a diffusion con trolled reaction. As a consequence, K21 must be larger than 103. On the other hand, a too large value of K21 would imply an unobserved large value of [IOOH] and the value of K21 is probably between 103 and 104.

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