Lab 3 - Chemical Kinetics: The Iodine Clock Reaction*
April 18, 2000

*Adapted from "Laboratory Experiments for Chemistry - The Central Science", 7th ed., J.H. Nelson and K. C. Kemp, Prentice-Hall, 1997.
 

Introduction. In this experiment you will determine the rate law for the following reaction:

In this redox reaction (Which species is being oxidized? Which is being reduced? What is the reducing agent?), the formation of water and iodine follows the rate law, k[H₂O₂]^m[I^-]ⁿ. The purpose of the experiment is to determine the order of the reaction with respect to each reactant, i.e., determine the values of m and n. From this information you will then calculate the value of the rate constant and propose a mechanism for the reaction.

To do any kinetics experiment you need to determine the rate at which species are consumed or generated. Often this is done by determining concentration data via absorbance or pressure measurements. In this experiment you will use a chemical indicator that provides a visible signal when the reaction has progressed a certain amount, that is, when a specified amount of reactants have been consumed. Starch forms a complex with iodine that is a deep blue color. Thus, if starch is present in the reaction system, the color of the solution will indicate the presence of molecular iodine. Therefore, if starch is present when the above reaction is performed the solution should immediately turn blue since iodine is generated directly. Obviously, if performed this way, no kinetic information can be gleaned since the color change would do nothing more than show that the reaction had started.

To extract rate data an additional compound is added to the reaction mixture. Thiosulfate reacts with molecular iodine very rapidly to regenerate I^- and the new compound S₄O₆^2- (see below). By adding a known amount of thiosulfate, the color change will then signal the point at which the thiosulfate has been consumed, provided that the amount of thiosulfate added is small relative to the amounts of peroxide and iodide initially present. This is because, at that time of the color change, the I₂ that is generated by the first reaction no longer is "scavenged" by the thiosulfate and can build up in concentration to permit the formation of the blue I₂/starch complex. Since the stoichiometry of both reactions is known, the amount of thiosulfate reacted can be correlated to the amount of peroxide consumed in a given time.

To obtain the rate law, the reaction rate (molarity/sec) will be measured at various concentrations of I^- and H₂O₂. The reaction rate can be obtained by a plotting the moles of H₂O₂ reacted vs. time. Each experiment will consist of multiple additions of thiosulfate; for each addition, the cumulative time is plotted on the x axis and the cumulative amount of peroxide reacted is plotted on the y axis. The resulting slope will have units of mol/sec. By dividing the slope by the volume will give the reaction rate in the desired units, M/sec. The rate law can then be extracted by determining how the reaction rate varies with different concentrations of S₂O₃^2- and I^-.
 

Procedure.

Part 1. Solution Preparation. Prepare the following four reaction solutions. Do NOT prepare them all at once; make them one at a time, proceeding to Part 2 before preparing the next solution.
 

	Solution A	Solution B	Solution C	Solution D
0.25 M KI solution	25.0 mL	25.0 mL	50.0 mL	12.5 mL
1% starch solution	1.0 mL	1.0 mL	1.0 mL	1.0 mL
0.1 M Na2S2O3 solution	1.0 mL	1.0 mL	2.0 mL	1.0 mL
Buffer	5.0 mL	5.0 mL	5.0 mL	5.0 mL
Deionized Water	63.0 mL	58.0 mL	37.0 mL	78.0 mL
Total Volume	95.0 mL	90.0 mL	95.0 mL	97.5 mL

Part 2. Collection of Rate Data. The reaction will proceed quickly upon mixing of the reagents, so be prepared to start timing immediately. To Solution A, add 5.0 mL of 3% H₂O₂; this can be measured in a graduated cylinder and added all at once. Begin timing as soon as the solutions "meet"; swirl the solution to mix thoroughly. Record the time at which the color changes, but do not stop the timer. Quickly add a 1.0 mL aliquot of Na₂S₂O₃. This can be delivered via a buret. The color should disappear, but eventually return to the dark blue. Record the time at which the color returns. Repeat the cycle until five aliquots of  Na₂S₂O₃ have been added.

Repeat this procedure with solutions B, C and D, each time adding enough hydrogen peroxide to bring the total reaction volume to 100.0 mL. The only variation on this procedure comes in the case of solution C, in which case you need to add 2.0 mL aliquots of the thiosulfate solution during the trial.

Analysis of Data. Prepare four plots of your data as described in the introduction, i.e., moles of H₂O₂ reacted vs. time and determine the slopes. Assume that the concentration of hydrogen peroxide is a weight/volume percentage, that is 3.0 grams of H₂O₂ per 100 mL of solution. Make a table of peroxide concentration, iodide concentration and the observed reaction rate to determine the rate law. Based on your data, calculate the value of k that each data set yield and report the mean and standard deviation. From this rate law, propose a mechanism that explains the reaction products and the rate law.