H. F. Shurvell Queen's University Kingston, Canada
The Kinetics of an Ionic Reaction A physical chemistry experiment
In this experiment, the reaction between hydrogen peroxide, iodide ions, and hydrogen ions is studied. By means of a very simple procedure, the reaction order with respect to hydrogen peroxide is determined and a rate constant obtained. The procedure can be modified easily to give the order with respect to iodide ion and other kinetic data, such as the rate constants involved in the overall rate equation. These additional experiments will be outlined after the basic experiment has been described. The first successful quantitative investigation in chemical kinetics was carried out on the reaction between hydrogen peroxide, iodide ions, and hydrogen ions, by Harcourt and Esson in 1867 (1). Many workers have since studied the reaction and a review of the early work has been published (2). More recently the reaction has been studied bpF. Bell and co-workers (3). An experiment based on the reaction has been described previously in THIS JOURNAL (4, and it appears often in textbooks and monographs on reaction kinetics (5). We should like to describe a simple modification of the reaction which has been used successfully for many years a t Queen's University for elementary courses in physical chemistry. The stoichiometric reaction between hydrogen peroxide and iodide ions in the presence of hydrogen ions is:
+
1 3 ~ 0 ~21-
+ 2H+ = 2H2O + 1%
The rate law consists of two independent terms which suggests that the reaction can proceed by two independent pathways (5).
If large excesses of hydrogen ion and iodide ion are initially present so that the concentrations of these ions remain essentially constant throughout the reaction, then eqn. (1) becomes:
d = k, [HzO,l dt where k,
=
k, [I-]
+ k, [H+] [I-]
tion from linearity is expected towards the end of an experiment. The course of the reaction is followed very simply by allowing the iodine liberated to react with a measured quantity of standard sodium thiosulphate solution. The thiosulphate has the effect of regenerating the iodide ion according to the reaction: 2BOsP- b = 21- + Sloe2Since this reaction is fast, no iodine will be liberated until the initial quantity of thiosulphate has been consumed. At this point, because starch solution is added to the initial reaction mixture, the solution will turn blue. By making further additions of thiosulphate and observing successive times of appearance of the blue color, the rate of the reaction with respect to the hydrogen peroxide concentration can be easily followed. It should be noted here that while the hydrogen ion and iodide ion concentrations are only reduced slightly by a dilution effect during the reaction, the hydrogen ion concentration is also reduced by participation in the reaction.
+
The Experiment
The student is provided with 3% HZO,, 1M KI, 3 M H2S04,and 0.1 M Na2S203. A standard 0.02 M KMnOp solution is also provided. A starch solution is freshly prepared by making a slurry of about 10 g of soluble starch with 2C-30 ml of cold distilled water and pouring this slurry slowly with stirring, into about 100 ml of boiling or nearly boiling water. A clear solution should result. A solution of approximately 0.05 M H202is prepared by diluting 30 ml of 3y0 HzOzto 500 ml with distilled water. This solution is standardized by pipetting 25 ml into a 250 ml Erlenmeyer flask, adding a few ml of 3 M BzSOa and titrating with the standard KMnOa solution until the first permanent pink coloration is observed. The equation for this standardization reaction is: 5H201
(2)
Equation (2) can be tested by measuring [H20z]at various times and plotting t versus log [H202]. A linear plot will indicate that the reaction is first order with respect to hydrogen peroxide and the slope of the graph will give a value for the apparent rate constant k,. However, because the concentrations of hydrogen ion and iodide ion do not remain constant,' a slight devia-
' I n fact the concentrat.ion of hydrogen ion diminishes more rapidly than t,he iodide ion, as will be seen lrtter.
+ 2Mn01- + 6H+
-
ZMn++
+ 50>+ 8Hz0
It should be noted that the 0.05 M H20zmust be freshly prepared and standardized before each experiment, because dilute solutions of hydrogen peroxide decompose on keeping. 150 ml of distilled water are placed in a 500 ml Erlenmeyer flask together with 20 ml of 1M KI, 10 ml of 3 M HzSOn,2 ml of 0.1 M Na&Oa, and 5 ml of starch solution. These solutions are rnjxed by shaking and allowed to reach room temperature, which should be noted. Then, using a fast delivery pipet, 20 ml of the 0.05 A4 H,O, solution are added, starting a stopwatch at Volume 44, Number 7 0, October 1967
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the start of the addition. The time at which the solution turns hlue is recorded and a 2 ml aliquot of Na?SpOa is added immediately. The time is recorded when the solution turns blue again and the procedure is repeated until about 10 additions of N&S,Oa have been made. After the first few additions, the time intervals between the appearance of the hlue color will become long and further additions of N&SpOsshould be reduced to 1 ml aliquots. From the initial volumes used and the amounts of N&S20aadded, the concentrations of H202remaining a t each time may he calculated. The results are plotted in the way described earlier, to show that the reaction is first order with respect to hydrogen peroxide. From the slope of the line, the rate constant kl is calculated. Results and Discussion
Details of a typical experiment are given in the table and a plot o f t versus log [H202]for this experiment is shown in Figure 1. The linearity of the plot indicates that the reaction is first order with respect to H202.
Details of a T v ~ i c o Ex~eriment l
Initial mixture
ml
Temperature k.
24T 4.26 X 1 0 - 8
.w-l
least squares computer program to give the best straight line of the form: y = Ax B. The value of kl given in the table was obtained from this calculation. The changes in [H+]and [I-] during the experiment may be easily calculated. For example, initially the concentration of iodide ion was 0.020 moles in 207 ml, that is 0.097 moles/l. After the addition of a further 11 ml of 0.1 M N&SpOathe concentration was reduced to 0.092 moles/l. This represents a reduction of 5%. On the other hand, one mole of hydrogen ions is consumed for each mole of Na?SZ03reacted. The initial concentration of H + was 0.290 moles/l and the final concentration 0.269 moles/l, which represents a 7y0 decrease in hydrogen ion concentration. Thus a slight deviation from linearity of the first order plot towards the end of an experiment is expected. Depending on the ambient temperature and the concentration of the diluted HpOz,the reaction may proceed too rapidly initially, or may terminate after only a few additions of N&S303. In such cases, adjustments of the initial amounts of KI, H202,and N&SpOaused may be made to give a more satisfactory set of data.
+
Further Experiments A plot o f t versus log [H20z] b r
o typical experiment.
It is perhaps instructive to give sample calculations for some points in Figure 1. The molarity of the diluted H,O, solution was found to be 0.0403M in this experiment, thus 20 ml contained 8.06 X l,0W4 moles. The total initial volume of reaction mixture was 207 ml. Hence the initial H,02 concentration was 3.89 X lo-' M. We may calculate the concentration of H202 remaining at the appearance of the first blue color as follows: the initial reaction mixture contained 2 ml of 0.1 M N&SzOa(i.e., 2 X lo-' moles). Noting that 2 moles of N&S,Oa are equivalent to 1 mole of H202,we see that 1 X moles of HP02have reacted a t the appearance of the first blue color. The volume of thereaction mixture remains a t 207 ml, hence the concentration of H202has been reduced to':
At the appearance of the first blue color, 2 ml of thiosnlphate were added so a t the second appearance of the blue color a further 1.00 X lo-' moles of HpOphave reacted and the volume has been increased to 209 ml. Hence the concentration of H202remaining is:
The data for the experiment was fitted by means of a 578 / Journal o f Chemical Edumfion
The reaction order with respect to iodide ion could be determined by a procedure similar to that described above. In this case the hydrogen ion and hydrogen peroxide concentrations should be initially in large excess. Equation (1) then becomes:
The constants kpand ka in eqn. (1) can be deduced if k, is obtained for several reaction mixtures, differing initially only in acid concentration (6). Then a plot of kl versus [H+], when divided by [I-], should give a straight line of slope ka and intercept k2. In this modification of the experiment, care must be taken to ensure that the initial concentrations of H202 and iodide ion are constant. This is achieved by subtracting the volume of 3 M H2S04used from the volume of water added to the initial reaction mixture, so that the total volume is constant. All runs must, of course, be camed out a t the same temperature. This experiment has been used successfully in our third year Physical Chemistry course and rate constants in good agreement with Reference (2) have been obtained. The effect of ionic strength could be studied by adding non-reacting salts, such as potassium chloride (3) or sodium perchlorate (6),to the reaction mixture.
Literature Cited (1) HARCOURT, A. G . V., AND ESSON,W., Phil. Trans. Roy. Soc., 157,117 (1867). (2) LIEBHAFSKY, H. A., AND MOHM~MAD, A,, J. Am. Chem. Soc., 55, 3977 (1933). D., AND WYNNE-JONES, W. F. (3) BELL,F., GILL,R., HOLDEN, K., J . Phys. Chem., 55,874 (1951).
(4) MCALPINE, R. K., J. CHEM.EDUC.,22,387 (1945). (5) See for example: BENSON,5. W., "The Foundations of
Chemical Kinetics," McGraw-Hill Book Co., New York, 1960, pp. 23 and 33; KING,E. L.,"HOWChemical Reactions Occur," W. A. Benjsmin, Ino., New York, 1964, pp. 80-82. (6) LIEBAAFSKY, H. A,, 38, 857 (1934).
AND
MOHM~MAD, A., J . P h w Chem.
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