Clock Reaction

The iodide/hydrogen peroxide clock reaction

The iodide/hydrogen peroxide clock reaction

Introduction

Several studies of the Mo(VI) and W(VI) catalyzed reaction of H2O2 with various substrates have been reported [1, 2, 3, 4, 5, 6, 7 and 8]. One of the first detailed studies was that of the oxidation of I- with H2O2 in the presence of molybdate ion in an acid solution. This study by Garcia and Lara [1] found that, in 0.08 M H2SO4 solution, the rate law for the reaction is:

rate=k[MoO4 2-]0.5[H2O2]0.5[I-]1.2

Later, Smith and Kilford [2] studied the same reaction with I- and H+ in excess over H2O2. They found that under these conditions the reaction was neither first nor second order in H2O2. A slight dependence on H+ was found in the range 0.025-0.20 M. They concluded that their results could best be explained by the sequence:

H2MoO4+H2O2 H2MoO4(H2O2) (rapid, K1)

H2MoO4(H2O2)+H2O2 H2MoO4(H2O2)2 (rapid, K2)

H2MoO4(H2O2)2+I- products (rate determining, k1)

with k1=3.3×102 M-1 s-1. No activation parameters were given. Studies by Thompson and co-workers [6] have also shown that the formation of the proposed peroxo complexes have large formation constants.

Recently, it has been reported that the reaction between H2O2 and an excess of I- in the presence of molybdate in an acid solution follows the rate law:

rate=k2[Mo(VI)][H2O2]

with the formation of either a mono or diperoxo complex being the rate determining step which is followed by the rapid oxidation of I-.

We have decided to investigate this apparent contradiction by employing an iodide clock technique to measure the true initial rate of reaction. We have chosen to work in an acetate buffered (pH 4.5) medium where there is little interference by the acid promoted, uncatalyzed (by MoO4 2-) reaction [9 and 10].

The Iodine Clock Reaction

This is a classic surprise reaction. Two colorless solutions are mixed and nothing appears to happen for nearly one minute. The solution suddenly turns dark blue. The demo can be used at many levels ranging from very introductory "gee whiz" to a p-chem discussion of competing reactions.

Materials

100 mL of Solution A (0.20 M H2SO4, 0.088 M H2O2)

100 mL of Solution B (0.0016 M Na2S2O3, 0.052 M KI, starch)

500 mL flask or beaker

Presentation

After appropriate introduction, mix the two solutions and swirl or stir. If appropriate, show consternation that nothing is happening. Your behavior depends on the degree of showmanship you wish to display. If you time the reaction, you can show the effect of temperature by mixing cold or warm solutions. The effect of concentration on rate can be shown by using solutions of greater or smaller concentration.

DISCUSSION

In the iodine clock reaction, there are really two processes happening simultaneously. The first is a slow reaction producing iodine:

H2O2(aq) + 2 I-(aq) I2(s) + 2 OH-(aq)

However, the iodine is never seen because of the second very fast process in which it is immediately reduced back to colorless iodide ion:

I2(s) + 2 S2O32-(aq) 2 I-(aq) + S4O62-(aq)

Thus, iodine is slowly formed and then instantly converted back to iodide ion until all the thiosulfate ion (S2O32-) is used up. At that point the iodine concentration shoots up ...