The “Chemical Kinetics” experiment was done to investigate the changes in the rate of reaction under the effect of concentration, temperature, and presence of a catalyst. It was determined that as the concentration of reactants and the temperature increases, the rate of the reaction increases as well. Also, the reaction was run by the presence of catalyst, and the rate of the reaction increased drastically in the presence of it. The order of the reaction with respect to each reactant was calculated to be: x = 1 [I-], y = 1 [BrO3-], z = 2 [H+] by the method of initial rates. The average rate constant was determined to be 26.7 M-3s-1, and the activation energy was calculated to be 49.6 kJ/mol.
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We notice that it happens as the temperature stays about constant. It can be noticed from trial 1 to trial 4. The calculated order for the reaction with respect to each reactant is as follows: [I-] was 1, [BrO3-] was 1, and the order for [H+] was 2, thus an overall order of 4.
Reaction 1 2 3 4 k (M-3s-1) 25.0 26.7 27.2 27.9 kavg. (M-3s-1) 26.7
Table 2 illustrates the numbers calculated for k and then the number for average k in M-3s-1. The calculation is done using equation (4).
I calculated the concentrations of [BrO3-] and [H+] for each mixture by using M1V1 = M2V2 equation. As can be noticed from Table 1, the rate of the reaction is really dependent on the concentration of the reactants. We notice that as the concentration of the reactants mentioned in the table increases, the rate increases as well. This means that the experiment was done right and it followed the rules of Kinetic’s theory; the