The Rate of Oxidation of Iodide Ion by Hydrogen Peroxide R. K. McALPINE University of Michigan, Ann Arbor, Michigan HE practice of chemistry is concerned very peroxide is present in excess, the sodium thiosulfate largely with the carrying out of chemical reactions will finally be used up and then the free iodine will whereby new substances are produced from old. The accumulate in the solution. In order to recognize the substances entering into the reaction are'known as free iodine while its concentration is still quite low, a reagents, and the substances produced as products. little starch solution is included in one of the original One of the factors in chemical reactions which is mixtures. Thus the solution remains colorless as long subject to human control and manipulation is sum- as any thiosulfate is present, but as soon as this is all marized in the Law of Mass Action. This law, in oxidized the solution turns blue. On the basis of earlier experiments a set of stock simplified form, states that a t a given temperature the rate of a reaction varies directly as the active mass of reagents was made up as follows: each reagent. Practically, the active mass of a re5N Sulfuric acid.. . . . . . . . . . . . . . . . . . .approximately agent is approximately identical with its concentration, 0.5 N Potassium iodide.. . . . . . . . . . . . . . .approximately Sodium thiosulfate.. .. .. . . . . . . . . .approximately 0 . 0 5 N especially in dilute solutions, although with higher Hydrogen peroxide (3 per cent concentrations additional factors known as activity 1.8 N solution). . . . . . . . . . . . . . . . . . . . .approximately coefficients may be introduced. Starch solution.. .. . . . . . . . . . . . . . .approximately 1 per cent Although this law is referred to early and late in the teaching of chemistry, there are relatively few exercises In later experiments 2 N HCl was used in place of the which can be placed in the hands of students or used sulfuric acid and Superoxol (30 per cent HzOz) was for lecture demonstration to show the experimental diluted and tried to see if the negative catalyst used to basis of this law. Some years ago it had been observed retard the decomposition of the ordinary hydrogen that the rate of oxidation of iodide ion by hydrogen peroxide might affect the results. No difference in peroxide in dilute acid solution is slow enough to results was obtained. An interesting experiment may be carried out to permit the development of a clock reaction of definite usefulness in this connection. This winter, for a course show the diierence in rates of the two principal rein general chemistry for some Navy students, the actions, oxidation of I- to Izby the H202,and reduction reaction was brought out, dusted off, and used. It of the Iz by the NaaSzOa. Place 50 ml. of distilled worked so well that i t was investigated somewhat water in a 250-ml. Erlenmeyer flask and add about 10 further, and the results are reported in the present ml. each of the acid and the potassium iodide solutions and 4 to 5 ml. of the starch solution. Fill a buret with paper. In the actual experiments two reactions are involved. the sodium thiosulfate solution. Now add about 5 ml. First, in acid solution hydrogen peroxide oxidizes of the hydrogen peroxide and swirl. The solution turns iodide ion to free iodine a t a moderate rate, the actual blue-black almost instantly. Next run in some of the rate being subject to control by modifying the con- sodium thiosulfate solution, with swirling, until the centrations of all three of the reagents. Second, free solution is decolorized, plus 1 to 2 ml. excess. W u l iodine is reduced very rapidly by sodium thiosulfate. to mix, and let stand. In a few seconds the color A third possible reaction, which might he expected to returns. Decolorize again with the sodium thiosulfate take place, namely, the direct oxidation of the sodium and vary the excess added. The blue-black color will thiosulfate by the hydrogen peroxide, actually takes return again but the time required will vary with the place so slowly that its rate is negligible as compared actual excess of thiosulfate added. This process may be repeated a number of times. This experiment leads with the other two. Thus it is possible to place a solution containing to the conclusion that the oxidation of the iodide ion sodium thiosulfate and potassium iodide in one con- is taking place a t only a moderate rate, but the retainer and a solution of hydrogen peroxide and sulfuric duction of the free iodine proceeds very rapidly. One of the early experiments tried was designed to acid (or hydrochloric acid) in another. The two are then poured together and mixed, whereupon the show that the time required for development of the hydrogen peroxide starts oxidizing the iodide ion to blue color varied directly with the amount of thiofree iodine. But the sodium thiosulfate reduces the sulfate used. For this, a standard oxidizing mixture free iodine back to iodide ion as fast as it is formed. was prepared by taking 100 ml. each of hydrogen If the quantities are adjusted so that the hydrogen peroxide and sulfuric acid, diluting to a liter, and
T
mixing thoroughly. Then four diierent concentrations of sodium thiosulfate were prepared by measuring out 5, 10, 15, and 20 ml. of the stock solution, treating each with 10 ml. of the potassium iodide solution and 5 ml. of the starch solution, and dilut'ng to 100 ml., followed by thorough mixing. In carrying out a reaction 25 ml. of the oxidizing mixture was placed in a clean 8-inch test tube and 25 ml. of the NazSzOa-KI-starchsolution was placed in a second test tube, using pipets for these (and earlier) measurements. A tenth-second stop watch was then started, the two test tubes picked up, and at the instant the pointer reached the 30-second mark the solutions were poured together and mixed by pouring back and forth four times. The tube was then held against a white background and the watch stopped a t the first showing of the blue color. The time recorded is the time from the initial pouring together to the first appearance of color. Three trials were run for each concentration of sodium thiosulfate and the average taken. The individual data are recorded in Table 1.
duration of the reaction. Lumped together, they may be considered as a "timing correction" to be subtracted from the over-all times to get values actually proportional to the times of the reactions. With the concentrations of sodium thiosulfate varying in the ratio 1:2: 3: 4, it is a simple matter to see whether a constant can be subtracted from each of the recorded times (10.8, 20.9, 30.9, and 40.8) to give a closer approach to the simple ratio. In this particular case, if each recorded time is diminished by 0.8 second the ratio becomes 1 :2.01 :3.01 :4.00. In connection with this timing correction, which is in the nature of a catch-all for several items, it is possible to determine experimentally how much free iodine is required to produce a blue color of sufficient intensity to be easily recognized. Several tubes of distilled water may be set up with a little K I and starch solution in each, and then single drops of a very dilute solution of iodine (0.0054 N in the case tried) added and mixed, and the colors noted. One drop of the iodine solution, corresponding to 0.035 mg. It produced a faint color, and two drops, or 0.07 mg., a medium blue. Since the amount of sodium thioTABLE 1 sulfate oxidized in 10 seconds required approximately RBLA~O OW NCONCBNTIAZION (IF Na2S20sro TKYB 10. APPBARAWCB 0 s co,.oa 8 mg. of free iodine, the time that elapsed between the end of the reaction and the development of the blue Relolilra Time i n Szrondr Conrrnlrnlianr Ratio of color was 0.07/8 times 10, or approximately 0.1 second. of NatSz08 Id Trial 2nd Trio1 3rd Trinl Avcrona Avrrnrrr Further, the physiological reaction time from the appearance of the color to the stopping of the watch may be assessed as approximately 0.3 second. This item could be measured experimentally with such It is obvious that the ratio of average times is suffi- special equipment as is nsed in the laboratory in ciently close to the ratio of concentrations to justify psychology, but it has not seemed worth while to check the statement that the time required to produce the this matter. It does appear, however, that in the blue color varies approximately with the amount of reaction being investigated and under the experimental sodium thiosulfate used. In attempting to account conditions used, there is surprisingly little error involved for the deviation on grounds other than "experimental in taking the instant of initial pouring as identical with error," two factors in the situation may be examined. the actual first point in the time of the reaction. One of these involves the fact that as diierent amounts Before presenting a summary of the further experiof sodium thiosulfate are nsed correspondingly diierent ments carried out, in which the concentrations of amounts of hydrogen peroxide and sulfuric acid are hydrogen peroxide, iodide ion, and acid were all varied, used up, so that the concentrations of these reagents it will be interesting to examine the constancy of the change during the course of the reaction. This factor timing results. Fifty-two experiments have been run is reduced to a very minor item by the relative amounts in triplicate. By comparing the first run with the of reagents used. The actual concentrations of hydro- second, first with third, and second with the third, gen peroxide and sulfuric acid change very little in all three checks or comparisons of times are available for four cases. Further, the change would be proportional each run, or 156 for the 52 experiments. The variations to the amount of sodium thiosulfate used so the ratio of of these among themselves are summarized in Table 2. times would not be affected. The second factor involves the relation of times recorded to the times of the actual reactions. This factor is more difficult to deal with. It is obvious that the two solutions are not mixed instantly, therefore there is some brief interval of time between the start of pouring together and the theoretical start of the reaction. Further, a short time must elapse beyond the end of the reaction and the development of the blue color, and a further time is required to stop the Total checks.. ................................ 156 watch after the color is recognized. Certain of these Maximum variation.. ......................... 1 . 3 seconds items are approximately constant, independent of the Average vsriafioo.. ........................... 0 . 2 8 second
In setting up the various experiments it was found desirable to adjust the different concentrations so that the maximum time would not exceed one minute. This avoided two difficulties, namely, straying attention while waiting for the end point, and decreasing sharpness of the end point. At the other extreme, it was necessary to use 10 to 15 seconds in order to carry through the standard mixing and he prepared to stop the watch. Thus, in varying the concentration of iodide ion 5, 10, 15, and 20 ml. of the potassium iodide solution were each treated with 10 ml. of the sodium thiosulfate and 5 ml. of the starch solution and diluted to 100 ml. These were then reacted separately with a standard mixture of hydrogen peroxide and acid (100 ml. of each, diluted to a liter). For variation of hydrogen peroxide or acid, 5-, lo-, 15. and 20-ml. portions of one were treated with 10-ml. portions of the other and diluted to 100 ml. These were then reacted separately with a standard reducing mixture of sodium thiosulfate, potassium iodide, and starch. Variations of this last mixture were tried, all the way from 200 ml. of sodium thiosulfate plus 200 ml. of potassium iodide solution (plus 40 ml. of starch solution) diluted to one liter, down to 10 ml. each of sodium thiosulfate and potassium iodide solution. I n the later experiments the mixture adopted was 50 ml. of each of the solutions, diluted to one liter and mixed. To show the variation in the rate of the reaction with the concentration of the various reagents, the raw timing data were converted to rates by comparing the reciprocals of the times observed. The unit of chemical reaction involved in each case was the liberation of the amount of iodine necessary to use up a standard amount of sodium thiosulfate. The longer the time required, the slower the reaction was proceeding. Thus the reciprocals of the times were used to compare the rates. These data are summarized in Table 3. The different series referred to involved such things as shifting from sulfuric acid to hydrokhloric acid, use of Superoxol in place of ordinary hydrogen peroxide, and variation in the concentrations of the standard reducing mixture. These are not discussed further. TABLE 3 RAT=
Conernuationr studied Rca~cnl "".id
-
nnrros sou RAW T m n o DATA(Rnrs 1:2:3:4
-
1/11
Rdcr
varying amounts of sodium thiosulfate. This is sufficient to indicate that the rate of oxidation of iodide ion by hydrogen peroxide, in the acid range studied, varies essentially as the first power of the concentration of each of the two reagents. In the case of both sulfuric and hydrochloric acid the effect does not follow a simple relation. For the first two members of series a, b, and c in the case of sulfuric acid, and series a and b for hydrochloric acid the effect is approximately a square root relation. That is, doubling the concentration of the acid increases the rate by approximately the square root of 2 (1.414). However, a fourfold increase in the concentration of the acid produces definitely more than a twofold increase in rate, and in series c of the hydrochloric acid studies the concentrations were double those of series b and a still more marked deviation toward a first power relation is observed. The rates considered in the last paragraph are those obtained from the raw timing data. It is possible t o calculate a timing correction from the experimental data for hydrogen peroxide and potassium iodide on the assumption that the rate is strictly a first power effect in these two cases. Thus, by applying a correction of one second the three series for hydrogen peroxide become 1:2.00:2.96:4.06 1:1.96:2.88:4.00 1:1.98:3.02:3.94
In the case of potassium iodide a correction of two seconds in series a (using 5 N sulfuric acid and ordinary hydrogen peroxide) and 2.5 seconds for series b (using 2 N HC1 and diluted Superoxol) gives the following ratios: Seriesa 1:l 94:2.98:4.08 Series b 1:2.01:3.01:3.98
In conclusion, it should be noted' that there are several anomalies in the system under study, an anomaly being defined as an unexpected irregularity in chemical behavior. Thus it is not a t all obvious why hydrogen peroxide should oxidize iodide ion so much more slowly than iodine oxidizes sodium thiosulfate, and it is still more difficult to understand why hydrogen peroxide should have so little effect on such a strong reducing agent as sodium thiosulfate which can he oxidized so rapidly by iodine. Further, i t may he noted that the chemical equation for the net reaction being studied is written:
+
HIOl 21-
In these raw rates it will be observed that for the several series of experiments involving hydrogen peroxide and potassium iodide, the ratios are qualitatively the same as the relative times required for oxidizing
+ 2H+ = Is + 2H10
According to the law of mass action, if the reaction took place directly in accordance with the equation it would be expected that the rate would vary directly as the first power of the concentration of hydrogen peroxide (which it does), and as the second power of the concentrations of both iodide ion and hydrogen ion. The fact that it goes as the first power of the iodide ion rather than the square has been accounted for by a
mechanism involving two reactions, one slow and one fast, the fitst forming hypoiodous acid (HIO),
+ H C+ I -
H202
=
HI0
+ H20
and the second forming free iodine by the reaction of the hypoiodons acid with more iodide ion,
plex than the one just suggested. Some further work is being done along this line. SUMMARY
A clock reaction has been developed to show in greater detail than usual the experimental basis of the HI0 + H + + I - = I* H20 law of mass action. Only a fourfold variation in conThe slowness of the first reaction controls the over-all centrations was investigated except in the case of rate, and in this equation the iodide ion has a coefficient hydrogen ion, but the range can be increased conof one so first power effects are to be expected (or are siderably by proper adjustment of the various reaccounted for). However, it is to be noted that the agents. In the oxidation of iodide ion by hydrogen coefficient for hydrogen ion is also one and so fust peroxide in moderately acid solution, the rate varies power effects would be expected in this case also. directly as the concentrations of the hydrogen peroxide The fact that the data for variation of hydrogen ion and the iodide ion, and less than directly as the concendo not show first power effects indicates that under the tration of the hydrogen ion. The effect of ionic conditions studied the mechanism must be more com- strength has not been studied.
+
