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.
The whole purpose of this experiment is to deal with the laws of chemical kinetics, and by doing the experiment, compare the experimental results with the theories and see if they were followed. From the kinetics studies, it is obvious that the rate of a reaction increases as the temperature of the reaction increases and as the concentration of the reactants increases. Also, the catalyst increases the rate of the reaction and decreases the activation energy. Thus, this experiment is divided into three sections and the dependence of the reaction rate from different factors is observed step by step. The idea of the first part of this experiment is to find the reaction orders with respect to each reactant and the rate constant, k. The method of initial rates to calculate the order with respect to each reagent will be used. With different times, the concentrations will vary in each trial. In this experiment, the following type of reaction is considered: aA + bB → cC + dD (1)
The rate law is:
Rate = k[A]x[B]y (2)
The reaction that occurs between the iodine ion and bromate ion will be studied, which is as follows: 6I-(aq) + BrO3-(aq) + 6H+(aq) > 3I2(aq) + Br-(aq) + 3H2O(1) (3) Then, the rate of the reaction is concluded to be:
Rate = k[I-]x[BrO3-]y[H+]z (4)
Since the rate of the reaction is obvious, the values for x, y, and z will be solved and afterwards, the rate constant, k, will be calculated. In the second part of the experiment, the dependence of reaction rate on temperature will be investigated in order to determine the activation energy, Ea, for the reaction. The rate of the reaction will be recorded at three different temperatures. In one case the reaction is going to happen in a hot bath and in another case, in a ice bath. Then the speed difference of the rates will be determined. Afterwards, the Arrhenius equation will be used to calculate the activation energy. The Arrhenius equation is as follows:
Finally, the reaction will be run in the presence of a catalyst and the dependence of reaction rate on a catalyst will be determined; whether it speeds up the reaction or slows it down. Experimental
The first part of the experiment states what mixtures to mix with each other in order to figure out the dependence of reaction rate on concentration. Three reagents (KI, Na2S2O3, H2O) were mixed into one flask (Reaction Flask 1), and other two reactants (KBrO3, HC1) into the second flask (Reaction Flask 2). The following table shows the amount of each reagent that is supposed to be mixed in each flask for each trial:
Reaction Flask 1(250 mL)Reaction Flask 2 (125mL)
110 mL10 mL10 mL10 mL10 mL3 drops
220 mL10 mL0 mL10 mL10 mL3 drops
310 mL10 mL0 mL20 mL10 mL3 drops
410 mL10 mL0 mL10 mL20 mL3 drops
58 mL10 mL12 mL5 mL15 mL3 drops
After the reagents for each flask were mixed with each other, the content of flask two was poured in flask one. The solution was swirled until the mixture turned blue. As soon the solution turned blue, the stopwatch was stopped and the time was recorded. The temperature was also recorded. The same process was done for the other four...