The Rate of Oxidation of Iodide Ion by Hydrogen Peroxide

T HE practice of chemistry is concerned very largely with the carrying out of chemical reactions whereby new substances are produced from old. The substances entering into the reaction are'known as reagents, and the substances produced as products. One of the factors in chemical reactions which is subject to human control and manipulation is summarized in the Law of Mass Action. This law, in simplified form, states that a t a given temperature the rate of a reaction varies directly as the active mass of each reagent. Practically, the active mass of a reagent is approximately identical with its concentration, especially in dilute solutions, although with higher concentrations additional factors known as activity coefficients may be introduced. Although this law is referred to early and late in the teaching of chemistry, there are relatively few exercises which can be placed in the hands of students or used for lecture demonstration to show the experimental basis of this law. Some years ago it had been observed that the rate of oxidation of iodide ion by hydrogen peroxide in dilute acid solution is slow enough to permit the development of a clock reaction of definite usefulness in this connection. This winter, for a course in general chemistry for some Navy students, the reaction was brought out, dusted off, and used. It worked so well that i t was investigated somewhat further, and the results are reported in the present paper. In the actual experiments two reactions are involved. First, in acid solution hydrogen peroxide oxidizes iodide ion to free iodine a t a moderate rate, the actual rate being subject to control by modifying the concentrations of all three of the reagents. Second, free iodine is reduced very rapidly by sodium thiosulfate. A third possible reaction, which might he expected to take place, namely, the direct oxidation of the sodium thiosulfate by the hydrogen peroxide, actually takes place so slowly that its rate is negligible as compared with the other two. Thus it is possible to place a solution containing sodium thiosulfate and potassium iodide in one container and a solution of hydrogen peroxide and sulfuric acid (or hydrochloric acid) in another. The two are then poured together and mixed, whereupon the hydrogen peroxide starts oxidizing the iodide ion to free iodine. But the sodium thiosulfate reduces the free iodine back to iodide ion as fast as it is formed. If the quantities are adjusted so that the hydrogen peroxide is present in excess, the sodium thiosulfate will finally be used up and then the free iodine will accumulate in the solution. In order to recognize the free iodine while its concentration is still quite low, a little starch solution is included in one of the original mixtures. Thus the solution remains colorless as long as any thiosulfate is present, but as soon as this is all oxidized the solution turns blue. On the basis of earlier experiments a set of stock reagents was made up as follows: Sulfuric acid.. . . . . . . . . . . . . . . . . . .approximately 5 N Potassium iodide.. . . . . . . . . . . . . . .approximately 0 . 5 N Sodium thiosulfate.. . . . . . . . . . . . . .approximately 0.05 N Hydrogen peroxide (3 per cent solution). . . . . . . . . . . . . . . . . . . . .approximately 1.8 N Starch solution.. . . . . . . . . . . . . . . . .approximately 1 per cent