Experiment X
Kinetics of the Decomposition of Hydrogen Peroxide

Name: Will Milliken Student Number: 6771898 Partner’s Name: Chris Cariou Course Code: CHM 2330 Due Date: January 28, 2014
Table of Contents
Objective………………………………………………………………………………………………………………………………..3
Introduction……………………………………………………………………………………………………………………………3
Procedure………………………………………………………………………………………………………………………….……5
Table 1…………………………………………………………………………………………………………………………6 Table 2…………………………………………………………………………………………………………………………6 Table 3…………………………………………………………………………………………………………………………7
Results…………………………………………………………………………………………………………………………………….8
Table 4…………………………………………………………………………………………………………………………8
Table 5…………………………………………………………………………………………………………………………8
Table 6…………………………………………………………………………………………………………………………9
Table 7…………………………………………………………………………………………………………………………9
Table 8…………………………………………………………………………………………………………………………10 Table 9…………………………………………………………………………………………………………………………10
Discussion………………………………………………………………………………………………………………………………..11
Table 10.………………………………………………………………………………………………………………………13 Table 11……………………………………………………………………………………………………………………….13
Conclusion……………………………………………………………………………………………………………………………….13
References…………………………………………………………………………………………………………………………….…14

Objective The goal of this experiment is to explore how varying concentrations and temperatures affect the rate of formation of iodine. The rate will be used to determine the rate constant, order of reaction, and activation energies of the formation of iodine. In addition, a catalyst will be used, and its effect on the rate will be determined from the experimental data.

Introduction In this experiment, the rate of reaction of iodine will be explored. Through the use of a buffer solution of Na2S2O3, as well as diluted solutions of KI, MoO¬4-2 and H2O2, the rate of iodine formation will be timed and compared to that of other trials in order to understand how different concentrations of the reactants affect the rate. This type of experiment is commonly known as an iodine clock reaction experiment. It is very important to understand what factors have an effect on the rate of a reaction for many reasons. Some real-world examples would be a company looking to slow down a reaction rate in order to prevent product spoilage (such as foods and chemicals), while others may be looking to increase the rate of the reaction, such as companies that produce digestive pills that react in the stomach. Without experiments such as the iodine clock reaction experiment, it would be hard to judge what concentrations of chemicals should be included in various products and solutions for the best results. The concentrations of each reactant will be calculated using the formula:
[reactant] = (M*Vreactant)/(VT) (1)

where [reactant] = concentration of reactant, M = molarity,
Vreactant = volume of reactant, VT = total volume
In order to understand how varying concentrations and catalysts affect the rate of reaction, it is first necessary to understand how a reaction happens. The rate of reaction is governed by the collision theory. Molecules collide in the solution and favourable molecules will react upon collision if the collision has enough energy to break/form bonds among the molecules.5 By increasing the concentrations of some of the reacts, the amount of favourable collisions will increase and therefore the rate of reaction should theoretically increase as well. In the same way of thinking, by decreasing the concentrations of other reactants that are not necessary for the desired reaction, this will make more room for the favourable collisions to occur, and thus increase the rate. Since the initial concentration of KI is directly related to the final concentration of I2, the initial rate of formation of I2 can be determined using the formula below:
Rate of Formation of I2 = [KI]/time