# Iodine Clock Reaction Primary Salt Effect in the Kinetics of Ionic Reactions Essay

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ABSTRACT: The purpose of this experiment was to determine the rate constants and ionic strengths of the series and to prove the Bronsted relation. In order to do so, known amounts of KI, Na2S2O3, KNO3, EDTA, starch and K2S2O8 were mixed up, and placed in the spectrophotometer until the %T reached 60%, and time was recorded. In the first part of the calculations, for flask 1, 2 and 3, the true reaction rate was calculated using the equation k = (1/∆t) x ([S2O32-]/[Iodine][S2O82-]). Which resulted in values of 2.8878765.66 x 10-3 s-1 , 3.159845 x 10-3 s-1, and 3.079703 x 10-3 s-1, these values are all similar to each other since they contain no electrolyte reacting with the persulfate solution. The apparent reaction rate was calculated using the equation, kapp= (1/∆t) x ([S2O32-]/[S2O82-]) which resulted in apparent rate constants of 5.66 x 10-5 s-1, 6.1958 x 10-5 s-1, 6.0356 x 10-5 s-1. The actual concentration was calculated using the basic chemical equation, C1V1 = C2V2. In order to find the order of reaction a a graph of log rate vs. log [S2O82-] was drawn, and was found that the results gave a zero order reaction But in reality the reaction order in [I-] and [S2O82-] is in first order each, although [I-] is kept at a constant volume throughout the reaction therefore the overall reaction is pseudo- first order. -d [S2O82-] = kapp [S2O82-] Dt In the second part the rate constant was found using the equation k = (1/∆t) x ([S2O32- ]/[Iodine][S2O82-]). Where it resulted to values of 3.990602 x 10-3 s-1, 4.653278 x 10-3 s-1, 5.944044 x 10-3 s-1, 7.499958 x 10-3 s-1, 7.499958 x 10-3 s-1, 9.84554 x 10-3 s-1, for flasks 4, 5, 6, 7, 8. Then the Ionic strength was calculated using the equation µ=½∑ ci zi2 . Which resulted in 4.5392142 x 10-2, 6.4999998 x 10-2, 1.23823528 x 10-1, 2.21862742 x 10-1, and 4.17941175 x 10-1. It