P. C. Moews, Jr. and R. H. Petrucci Western Reserve University
Cleveland, Ohio
The Oxidation of Iodide Ion by Persulfate Ion
In a recent symposium (1) sponsored by the Division of Chemical Education, ACS, the need for further reactions amenable to simple kinetic studies was emphasized. More specifically it was stated (3) that reactions are needed whose rates can be measured readily, for which the rate equation can be determined directly, and for which it is obvious that the rate e q u a tion and the equation for the net reaction are not related in any simple way. A kinetic experiment carried out by large sections of freshmen a t Western Reserve University-+ study of the rate of oxidation of iodide ion by persulfate ion-eems to fit most of these criteria. moreover the experiment provides for a study of the effect of temperature, influence of ionic strength, and effect of a catalyst on the rate of a chemical reaction. The reaction between persulfate ion and iodide ion may be represented by the equation:
The reaction is second order, the rate being dependent upon the concentration of both I- and S20s2-. Mack and France (3) employed this reaction in a physical chemistry experiment which clearly illustrates all of the
points mentioned in the previous paragraph. Mack and France's experiment, while excellent for undergraduate physical chemistry, takes more than one laboratory period and is somewhat involved for freshmen. Evans ((4 has suggested the persulfate ion-iodide ion reaction as being ideal for a kinetic experiment in general chemistry. Evans made use of an initial rate method in which solutions were allowed to react until a definite concentration of iodine, estimated visually, was attained. A more convenient initial rate method, first described by Icing and Jacobs (5),involves the reduction of an aliquot of sodium thiosulfate by the iodine formed in the reaction, 2&0,'-
+ 1,-
=
&OP
+ 31-
When all of the S20? is consumed, the liberated iodine can be detected with starch indicator. Two recent laboratory manuals (6, 7 ) describe kinetic experiments employing this reaction. I n one (6, Wilson, et al.) the reaction is followed by titration; in the other (7, Faigenbaum, el al.) an initial rate method similar to that of King and Jacobs (5) is used. Our experiment also utilizes Icing and Jacobs initial rate method, but
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Table 1.
ExperiMolarity ment number (NH&S208 1 2 3 4 5 a
0.077 0.038 0.019 0.077 0.077
The Effect of concentration of KI and
Molarity
(NHdZ~20s on the
Column A
(NH.)804
(seo)
k X 10' (1 mole-' seo-')
0:038 0.059
21 44 91 42 79
62 60 57 62 66
t
KI
KNOI
0.071 0.077 0.077 0.038 0.019
... . .. 0:038 0.059
... . ..
Reaction Rate (24
* 0.5°C)a
Bk X 10'
-Column
t (sec)
(1 mole-' sec-l)
Molarity KNO.
21 51 114 46 93
62 51 46 57 55
o:ii5 0.176 0.038 0.059
-Column
C-
k X 10" (1 mole-' (sec) sec-') t
21
62 62 63 62 66
42
81 42 79
All solutions contain 5 ml oi 0.2% starch and 10 ml of 0.01 M Na&Os, and have a total volume of 65 ml.
differs from the experiment of Faigenbaum, et al. in that reaction times are short, allowing several experiments to be performed during a brief laboratory period. The Experimenl
The student is furnished with 0.01 M N&Oa, 0.2% starch solution, and solid KI, KNOs, (NH4)zSzOs,and (NHJISOI. A stop watch is desirable but sweep second hand wall clocks may be substituted. The students prepare 0.2 M solutions of (NH&S20s and KI. From these solutions they prepare 0.1 M and 0.05 M solutions of both (NH&&Os and K I by dilution. I n order to keep the ionic strengths constant the students are told to use 0.2 M KNO3 and 0.2 M (NHJ2SOnto effect the dilutions of thc 0.2 M KI and 0.2 M (NHa)zS208, respectively. I n Experiment 1, 25 ml of 0.2 M XI, 5 ml of 0.2% starch solution, and a 10 ml aliquot of 0.01 M Na2S2O3 are mixed in a 250-ml beaker.' Twenty-five ml of 0.2 M (NHl)zS20sis quickly added to the beaker and the time observed simultaneously. Stirring is commenced immediately. At the appearance of a blue color the elapsed time is recorded and a temperature reading taken. Experiments 2-5 are carried out in the same way as Experiment 1 except that the concentrations of K I or (NH&320~ are varied. In Experiment 2, 25 ml of 0.1 M (NH4)2S208 (the solution is also 0.1 M in (NH& SO6) is substituted for the 25 1111 of 0.2 M (NHJ1S208 used in Experiment 1. Experiment 3 uses 25 ml of 0.05 M (NH4)zS20s; Experiment 4, 25 ml of 0.1 M K I ; and Experiment 5, 25 ml of 0.05 M KI. Some typical results are given in Table 1, Column A. The influence of ionic strength on the reaction can he shown by carrying out the dilution of the 0.2 M K I and (NHJ2S20s with distilled water and repeating the several experiments; Table 1,Column B gives some typical data. Some results obtained by maintaining the ionic strength of all solutions constant with KNOa are listed in Table 1, Column C. The effect of temperature is shown by heating and cooling the solutions before mixing them. The student is told to carry out Experiment 4 a t two temperatures below and a t one temperature above room temperature. Some typical results are listed in Table 2. Finally the effect of a catalyst on the reaction can be
' It may be necessary to add a drop of 0.1 M EDTA (ethyleu-
diaminetetraacetic acid) solution if the reagents or distilled water contain any appreciable quantity of metal ions which act as catalysts. Of course in such cases the addition of 0.1 M EDTA should he omitted when studying the catalytic effect of CUR+ions.
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Journal of Chemical Education
Table 2. The Effect of Temperature on the Reaction Raten
mint number
T ("C)
t (sec)
6 7
3
189
--
1.7
RR . .
24 33
42 21
8 9
(I mole-' sec-') 14 29 62 120 ~
Experiment 4, Tahle 1, Column A, repeated at four different temperatures.
studied by adding small amounts of CuZ+to the reaction mixture. The student is told to repeat Experiment 4 with the addition of one, two, and then three drops of 0.02 M Cu(NOa)z. Some typical results are given in Table 3. Interpretation of Student Data
Rate constants can be calculated readily from the data in Table 1. We have calculated rate constants using the equation
The rate constants appear in Tables 1 and 2. A value for the Arrhenius activation energy may be calculated from the data in Table 2 by plotting log k versus l / t (OK). A linear plot is obtained from which a value for the activation energy, 12.4 kcal, can be derived. For beginning students, however, we believe the determination of the order of the reaction to be a more important exercise than the calculation of rate constants. The student is told that one might expect the rate law for the reaction of I- with S2OS2-to be of the form: Rate
=
k[I-]"[S2OBP-1"
where the exponents m and n are small whole numbers. The student is asked to determine the values of the exponents m and n. The time that elapses from the mixing of the reactants to the appearance of the blue color is inversely proportional to the rate of the reaction. Plots are made of reaction time versus [I-] (with [S2082-] held constant) and reaction time versus [S20s2-] (with [I-] held constant). These plots are Table 3.
The Effect of a Catalyst on the Reaction Rate"
Experiment number
Drops 0.02 M Cu(N0a).
1
(see)
a Experiment 4, Table 1,~ o l & A, n repeated with the addition of various amounts of 0.02M Cu(NO.)..
used by the student to determine the dependence of the reaction rate on the concentrations of both I- and S20s2- and hence the values of m and n. Instead of a determination of a value for the activation energy, the profound effect of temperature on the rate of the reaction is impressed upon the student. Finally the student is asked to draw conclusions about the effect of a trace of Cu2+on the rate of the reaction from data like those in Table 3. Discussion
The reaction of S20a2-with I- has been investigated extensively; the first kinetic study is evidently that of Price (8) in 1898. Price concluded that the reaction is a bimolecnlar one and observed that copper and iron salts are effective catalysts. The catalytic effect of copper and iron salts has been confirmed by other workers (9). The reaction between persulfate ion and thiosulfate ion is also catalyzed by copper salts (10). This makes an exact interpretation of the data in Table 3 somewhat difficult; but this is unimportant for the purpose of showing the catalytic effect of trace amounts of copper. The reaction of persulfate ion with thiosulfate ion in the absence of a catalyst is insignificant since the rate of the uncatalyzed reaction is much slower than that of the S2Os2--I- reaction (10). I t has been suggested2that this experiment would be more effective if it were experimentally demonstrated that the reaction of S20a2-with S2032-does not occur to an appreciable extent. This was done by repeating the several experiments in Table 1, Column C, replacing the XI with KN03. After the solutions were allowed to stand for periods of time corresponding to those reported in Table 1, Colun~nC, they were diluted with 200 ml of cold water. A 0.005 M solution of 1% in 50% ethanol was used to titrate the unreacted thiosulfate. A typical set of results is: Experiments 1, 2, and 3, no loss of thiosulfate within experimental error; Experiment 4, 1% of thiosulfate consunled; Experiment 5, 3% of thiosulfate consumed. The most consistent results are observed if a drop of 0.1 M EDTA is added to each solution prior to the addition of the (NH4)2SzOs solution. An ionic reaction such as this which involves two negative ions has a rate which is strongly dependent upon ionic strength; it increases with increasing ionic strength. King and Jacobs (5) were able to interpret rate data a t low ionic strengths in terms of the BronstedDebye equation, log k = log k,
+ 2AZtZ9
Salt effects on the rates of ionic reactions a t higher ionic strengths are not well understood (11). Later workers have been concerned with the rates of the persulfate ion-iodide ion reaction a t higher ionic strengths and with the effects of temperature and solvent on the reaction. At higher ionic strengths specific effects have been reported (9) not interpretable in terms of the Bronsted-Debye equation or any of its modifications. Among the studies of the reaction of S20s2-with Ia t high ionic strengths is that of Meretoja (12). M e r e toja studied the effect on the reaction rate of ionic strengths varying from 0 to 2.5, KNOa being employed 2
Private communication, E. L. King.
as an inert electrolyte. He found that the effect of ionic strength could he represented with the semiempirical equation,
where ko is the velocity constant a t 11 = 0, A is the Dehye-Huckel constant, and or and B may he treated as parameters. I n view of more recent work this equation should not be expected to be generally valid. However it should be expected to he a t least approximately correct where the salt effects are largely those of 1:1 electrolytes. In fact using Meretoja's values for the above equation a t 24'C,
the calculated rate constants agree very well with those determined experimentally (Table 1, Columns A and C). The most consistent set of values for the rate constants (Table 1, Column C) are obtained when the ionic strength is kept constant with KNOa. Anunonium sulfate was used in Experiments 2 and 3 (Table 1, Column A) because it aUows salt effects to be explained to beginning students without the necessity of defining ionic strength. Conclusions
Using data for the persulfate ion-iodide ion reaction from the literature (e.g., that of Aleretoja) one can devise initial rate experiments to suit various purposes. Experiments can be designed to show the effects of temperature, catalysts, ionic strength, and reactant concentrations. The experiments can be devised to take varying lengths of time depending on the accuracy desired. The experiment as described in this paper is brief, and uses con~n~only available equipment. It is suitable for a short laboratory period and also has been very successful as a lecture demonstration experiment. While it fits most of the requirements for a simple undergraduate kinetic experiment the large salt effects do add a complicating factor. Literature Cited (1) KING,E. L.,Sympo~iumChairman, J. CHEM.EDUC.,40, 573-91 (1963). 40, 583 (1963). (2) CLVPBELL, J. A,, J. CHEM.EDUC., (3) EVANS,G. G., J. CAEM.EDUC.,29,139 (1952). W. G., "A Laboratory Manual (4) Mars, E. JR., AND FRANCE, of Elementary Physical Chemistry," 2nd ed., D. Van Nostrand Company, New York, 1934, p. 179. M. B., J . Am. Chem. Soc., 53, (5) KING,C. V., AND JACOBS, 1704 (1931). A. R., AND (6) WILSON,J. M., NEWCOMBE, R. J., DENARO, RICKETT, R. M. W., "Experiments in Physical Chemik try," Pergamon Preas, New York, 1962, p. 57. L. V., RICHTOL, H. H., AND (7) FAIGENRA~M, H. M., RACSTER, WIBERLEY, S. E., "Laboratory Manual for General Chemistry," John Wiley and Sons, New York, 1963, p.
--
14.
(8) PRICE,T. L., Z. physik. Chem., 27, 474 (1898). (9) HOUSE,D. A,, Chem. Rev., 62, 197 (1962). J. O., J . Am. Chem. Soc., 74, (101 . . SORUM.C. H., AND EDWARD. 1204(1952); (11) DAVIEG, C. W., "Progress in Resetion Kinetics," Pergamon PPMR. Ynrk. 1961. Vol. 1., D. 161. - - - - ~New -, (12) MERETOIA, A,, A m . Acad. Sci. Fennieae, Ser. A., 11. Chem. No. 24, 59 pp. (1947); C. A., 42, 2163c (1948).
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