2204
J . Phys. Chem. 1986, 90, 2204-2205
Iodine Chloride: A Key Intermediate in the Chlorite-Iodide Reaction Mihily T. Beck* and Gyula Ribai Department of Physical Chemistry, Kossuth Lajos University, Debrecen, 10, Hungary 4010 (Received: August 26, 1985)
The enigmatic oligo-oscillatorybehavior of the chlorite-iodide system, that is, the appearance of two extrema on the iodine concentration vs. time curves, can be explained by the formation and subsequent disproportionation of iodine chloride. The ratio of the iodine concentration consumed in the second phase to that regenerated in the third phase is 5 : 2 , as required by the proposed reactions. The characteristic shape of the kinetic curves is perfectly described by considering the rate equations of four-component reactions. The partial regeneration of iodine has been observed in the chlorite-iodine reaction, too.
The oxidation of iodide by chlorite may exhibit a number of most interesting kinetic phenomena.' In a thorough analysis of this reaction Epstein and Kustin2 mentioned that the batch reaction is in fact oligc-o~cillatory:~the concentration of iodine first reaches a maximum, then decreases to a minimum, and finally increases again until a final value is reached. Epstein and Kustin emphasized that this behavior cannot be explained in terms of the mechanisms suggested so far. The puzzling feature of the reaction is that when the concentration of iodine reaches the maximum value, the concentration of iodide decreases dramatically (below M), and remains at this level in the further stages of the reaction. Thus apparently there is no source of regeneration of iodine. We offer here strong evidence that for this behavior the formation of iodine chloride in the reaction between chlorite and iodine and its subsequent hydrolytic disproportionation is responsible. (The possibility of formation of IC1 has been considered and discarded by Grant et al.4 in the evaluation of the kinetics of chlorite oxidation of iodine.) Experimental Section A stock solution of 0.1 M iodine chloride was prepared by adding 1.100 g of potassium iodate dissolved in 5 cm3 of water to 0.71 3 g of potassium iodate dissolved in 9 cm3 of concentrated HC1. After the completion of reaction the volume was adjusted to 100 cm3. The mixing must be made in a closed vessel to avoid the loss of chlorine. The solution was 1.0 M for HC1. All the reagents were of analytical grade; iodine was dissolved in glacial acetic acid. Sodium chlorite stock solutions were freshly prepared from NaCIOz recrystallized from water; to prevent decomposition the solution was 0.001 M in sodium acetate. The pH of the reaction mixture was adjusted by acetate buffer. A Beckman Acta 111 spectrophotometer with a thermostated tandem cell was used for the experiments. The change in iodine concentration was monitored at 469 nm, the isosbectic point of 1, and 13-, The temperature was kept at 25 "C. The pH of the reaction mixtures did not change during the reaction. Results and Discussion Stoichiometric Considerations. Figure 1 shows the change in iodine concentration with time at constant initial iodide and varying initial chlorite concentrations. The ratio of the decrease in iodine concentration from the maximum to the minimum to the increase in concentration from the minimum to the final value was always found to be 5 2 , within experimental error. The change in concentration of iodine in the second and third phases of the reaction, that is, the decrease and the subsequent (1) Dateo, C. E.; Orbin, M.; De Kepper, P.; Epstein, I. R. J. Am. Chem.
increase in the iodine concentration as well as the observed ratio of these concentration changes, can be explained by the formation of iodine chloride in the reaction between chlorite and iodine and 21,
-
+3C10~
2 1 ~ 1 +2 1 0 ~ -+ CI-
by the partial regeneration of iodine in the well-known hydrolytic disproportionation of iodine chloride: 5IC1 + 3 H z 0
-
+ IO3- + 5C1- + 6HS
21,
Obviously, this explanation requires that in the reaction of chlorite with iodine the initial fast decrease of iodine concentration is followed by a slower increase until 215 of the initially consumed iodine is regenerated. This behavior was not observed by Grant et a1.: probably because of the relative slowness of the third phase and due to the fact that in most of their experiments a rather high excess of chlorite was used. Figure 2 shows, however, that in fact there is a minimum in the iodine concentration vs. time curves and the ratio of the originally consumed and regenerated iodine concentration is 5:2. It is evident from our experiments that the ratio A[12]:A[C10J depends on both the absolute and relative values of the concentration of iodine and chlorite and on the pH. This due to the fact that beside the chlorite oxidation of iodine some decomposition of chlorite also occurs in the solution, and the rates of these decomposition (disproportionation) reactions are influenced by the different species formed in the ClO,--iodine reaction. Kinetic Considerations. The two-extrema curves of Figure 1 can be semiquantitatively described by the rate equations of four-component reactions analogously with other oligo-oscillatory rea~tions.~*~Jj The first reaction is the oxidation of iodide by chlorite: 41-
+ C10; + 4H+
-
21,
+ C1- + 2 H z 0
(1)
The rate law and rate constants have been determined by Kern and Kim:'
k l a = 920 M-, s-1;
k l b = 5.3 x 10-3 s-I
Below M iodide the rate of this reaction was not taken into consideration in our calculations. We assume that the second reaction is the oxidation of iodine by chlorite: 21,
+ 3ClO2-
-+
2IC1
+ 2103- + C1-
(2)
The rather fast oxidation of iodine by chlorite has been quanti-
SOC.1982, 104, 504-509.
(2) Epstein, I. R.; Kustin, K. J. Phys. Chem. 1985, 89, 2275-2282. (3) Beck, M. T.; Ribai, Gy. J. Chem. SOC.,Dalton Trans. 1982, 1687-1 689. (4) Grant, J. L.; De Kepper, P.; Epstein, I. R.; Kustin, K.; Orbin, M. Inorg. Chem. 1982, 21, 2192-2196.
0022-3654/86/2090-2204$01.50/0
(5) Ribai, Gy.; Bazsa, Gy.; Beck, M. T. In?. J . Chem. Kine?. 1981, 13, 1277-1 288. (6) RBbai, Gy.; Beck, M. T. J. Chem. Soc., Dalton Trans. 1985, 1669. ( 7 ) Kern, D. M.; Kim, C.H. J . Am. Chem. SOC.1965, 87, 5309-5313.
0 1986 American Chemical Society
The Journal of Physical Chemistry, Vol. 90, No. 10, 1986 2205
Iodine Chloride
a
,,i
0
5
IO
b.
15
20
n
"0
2
L
8
6
10
12
I&
16
10
20
tJmin
Figure 1. Change in the iodine concentration with time for the iodidechlorite reaction. pH 3.40 (0.5 M acetate buffer), [I-lO= 2.0 X [CIOJ, = (a) 2.5 X (b) 2.0X lW3:(c) 1.5 X (d) 1.25 X 1W3
M.
Figure 3. Comparison of measured (-) and calculated (---) changes in the concentrationof iodine as a function of time for the iodidechlorite reaction. (a) [I-lO= = 2.0 X 10"; [C1OFl0 = 1.25X lo-) M;pH 3.40. (b) [I-lO= 4.0 X lo4; [C102-]o= 2.5 X lo4 M; pH 3.3. The experimental curve was obtained by Dateo et al.'
For our calculations is was sufficient to assume a bimolecular reaction and to use a rate constant larger than lo3 M-' s-l. A crucially important reaction is the hydrolytic disproportionation of iodine chloride: 5IC1 + 3 H 2 0
-
IO3- + 5C1-
+ 212 + 6H+
(4)
As far as we know, there has been no kinetic study of this reaction. According to our experiments the rate of the reaction is strictly proportional to the square of the iodine chloride concentration and strongly influenced by the p H of the solution:
-5 -w -2 1 - k4[ ICl]z 2 dt k4 =
-
O
0
8
ii
Figure 2. Change in the iodine concentration with time for the iodinechlorite reaction. pH 3.46 (0.5 M acetate buffer), [CIO,-]o = (a) 5 X lo4; (b) 1.0 X lom3;(c) 2.0 X lo-' M.
tatively studied by Grant et al.4 Under the conditions of oligooscillatory behavior they found the following rate equation:
We applied their experimentally determined value for k i but do not agree with their assumption that k - i = kf, since there is no evidence that the two constants refer to the same type of chemical change. All assumptions of this kind are rather arbitrary since the constants cannot be unambiguously assigned to a particular reaction. In our calculations the following constants were applied: k2/ = 0.54 s-l;
k-2/ = 4.5 X 10'O M-' s-l
kf = 2
X
lo7 M-I s-'
-
Iodine chloride reacts extremely fast with iodide: IC1
+ I-
I2
+ c1-
(3)
(4')
k4a 7 +- k4b [H 1 [H+I2
k4 = 40 M-' s-l at pH 3.40. The reaction is strongly inhibited by chloride. The value of k4 refers to zero chloride concentration, obtained by extrapolation. A detailed kinetic study of this rather complicated reaction will be published elsewhere. The reaction between iodate and iodide is negligibly slow under the conditions of oligo-oscillatorybehavior. Figure 3 shows that there is a qualitative agreement between the calculated and experimental curves. Figure 3a refers to our experiment, while the measured rate curve in Figure 3b is taken from the paper by Dateo et al.' By a fitting procedure a much better agreement could be achieved, but we do believe that a more quantitative treatment of data requires that much more information be available on the kinetics of the chlorite-iodine reaction, and on the kinetics of acidic decomposition of chlorite under the conditions of the oligo-oscillatory reaction. It is likely that the formation and reactions of iodine chloride play also role in the chlorite-iodide oscillatory reaction observed in a CSTR. Acknowledgment. A helpful discussion with Professors I. R. Epstein, K. Kustin, and M. Orban is gratefully acknowledged. Registry No. CIOT, 14998-27-7; I-, 20461-54-5;ICI, 7790-99-0.
