EXPERIMENT 12 KINETICS OF THE IODINE “CLOCK" REACTION PURPOSE The purpose of this experiment is to investigate the kinetics of the reaction of iodide and peroxydisulfate ions in order to determine the rate law of the reaction and the activation energy of both with and without a catalyst. METHOD The net equation for the overall reaction of iodide ion with peroxydisulfate ion is shown below: 2 (aq) + S2082(aq) → 12 (aq) + 2 SO4² (aq) (slow enough to measure) (1) (The 12 actually stays in solution as I3. by the reaction: I₂+ I → 13-) The rate law will be determined by the initial rate method. The concentration of each reactant will be systematically varied while holding the concentration of the other reactant constant. Each reaction studied will proceed until a fixed amount of thiosulfate ion has reacted with the product 12 from the first reaction. I2 (aq) + 2 S2032 (aq) → 2 I (aq) + S406² (aq) (fast) (2) S2082 (aq) + 2 S2032 (aq) → 2 SO4² (aq) + S406² (aq) (overall) (3) A starch indicator is added which, in the presence of 12, forms a starch iodine complex which is deep blue in color. As long as there is S2032 present in the solution, any I2 is immediately converted to I. However, when all of the thiosulfate ion is gone, the I2 builds up, and the solution changes from colorless to blue. The rate of reaction (1), R, is determined by measuring the rate of disappearance of thiosulfate ion in reaction (2), -d[S2032]/dt. Because each 12 consumes two S2032 ions, the rate of disappearance of S2O3 in reaction (2) is twice as fast as the rate of production of 12 in reaction (1). -d[S203]/2dt=+dt [12] /dt This allows calculations of R, the rate of reaction (1), as R=-d[S2O3²]/2dt = +dt [12] /dt = k [I] * [S203²¯] (4) (5) where k is the rate constant, x is the order of I, and y is the order of S₂O8². By comparing the values of R to the changes in initial concentrations of the two reactants, I and S2O8, the values of x, y, and k in rate law (5) can be determine following the initial rate method.