W. H. Cone and R. A. Hermens University of Idaho
Moscow
A Simple Kinetics Experiment For general chemistry laboratory
Simple experiments illustrating basic facts are needed in undergraduate chemistry. An experiment is proposed which requires no elaborate equipment. The time required for this experiment is about 45 minutes for preparation and three hours for following the reaction. Since potassium peroxodisulfate is a very strong oxidizing agent (oxidation potential = -2.01 v)' it can oxidize a variety of inorganic ions as well as certain organic compounds. I t has been found by Bacon and Doggart2 that, in the presence of catalytic amounts of silver ion, peroxodisulfate ion will oxidize benzoic acid to a resinous product. The method of measuring the rate of this reaction was by an iodometric titration of the remaining peroxodisulfate. Certain inorganic ions cannot be used since their oxidized state can be reduced by iodide ion which is used in the determination of peroxodisulfate ion. The resinous product formed from the reaction of benzoic acid does not interfere with the analytical determination. The reaction rate was found to be dependent on the concentration of two species, peroxodisulfate ion and silver ion. The rate was independent of the concentration of the reducing agent. The rate of the reaction can be expressed as follows:
The Experiment
In the absence of a catalyst the reaction rate of the peroxodisulfate ion is very slow. Allenahas found that uncatalyzed peroxodisulfate has a half-reaction time of 20 days a t 40°C in a reaction which is independent of the concentration of the reducing agent. A solution was made which was 0.01 11d' with respect to both benzoic acid and peroxodisulfate ion. To this was added enough silver nitrate solution to make the solution 0.001 M with respect to silver ion. The reaction mixture was placed in a constant temperature bath at 35°C. Aliquots were withdrawn and were determined iodometrically by the method of Kolthoff and Cam.& The liberated iodine was titrated with standard thiosnlfate using starch as an indicator. The figure shows a typical first-order plot of the data; k calculated from the slope is 1.41 min-', t1/2 = 492 min. Concentrations of catalyst between 0.0005 M and 0.0015 M may be used. Students' data should be within 5% error. "ALLEN, T. "OLTHOFF,
L., J . Am. Chem. Soc., 73, 3589 (1951). I. M., AND CARE,E. M., Anal. Chem., 2 5 , 298
(1953).
Since silver ion is present only as a catalyst its concentration remains constant in any particular reaction. The reaction, then, is first order with respect to the peroxodisulfate ion concentration. Then by integrating the result, equation (2) is obtained.
' LATIMER,W., M., "Oxidation Potentials," 2nd ed., PrenticeHall, Inc., New York, 1952, p. 78. 'BACON, R. G. R., A N D DOGGART, J. R., J . Chem. Sm.,1960, 1332-8.
Volume 40, Number 8, August
1963
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421
