Rate Laws: The Iodine Clock Reaction

Joseph Bronson

Lab TA: Longteng Tang
Lab Partner: Mychal Westendorf
Lab Section: Thursday 12:00-2:50
Abstract:
The purpose of this lab was to determine the rate constant, k, of an Iodine Clock Reaction based on the concentration at room temperature, and then constant concentrations at varying temperatures. This experiment proposes a relevance of a solutions temperature and concentration to how quickly the activation energy will be succumbed, which can be evaluated by determining the reaction’s rate constant. In part one of the Iodine Clock Reaction, the temperature was the constant variable. Using the formula , [I], [BrO3], and [H] were determined. After the concentrations were obtained, the natural log values of each of the concentrations were plotted against the natural log of the rate, and the orders of the reactants were determined to be [I]=1, [BrO3]=1, and [H]=2. Part two consisted of the same practice, except the amounts of each reactant were held constant for the entire experiment, and the temperatures were increased by increments of ten degrees in order to determine the rate constant as the particles began to move more quickly.

Theory:
In this laboratory, the rate constant needs to be determined in order to know how quickly the reactants of the Iodine Clock Reaction are being consumed, and in return, how quickly the product(s) may be forming. If a reaction is completed precisely and accurately, then the value of k should be the same for all concentration combinations at the specific temperature the experiment was conducted. Below are the chemical equations for this experiment:

According to the equation below, the rate of a reaction is equal to the rate constant, k, multiplied by the concentration of A raised to its order, multiplied by the concentration of B raised to its order. The value of k will always be positive, but if reactants are being consumed, a negative sign must be multiplied into the equation in order to account for the diminishing value that the rate will need to equal. The orders n and m must also be determined experimentally (and in cases where there are more than two reactants, two or more orders) in order to calculate k.

By taking the natural log of the rate and plotting it with the natural log of the concentrations of each of the reactants (I, BrO3, and H), a graph is created showing the linear regression of the reactants, where is used to determine the slope of each, and therefore the order of each reactant.
In reaction number one (refer to page one), Iodide and Bromate ions are reacting in an acidic solution to produce di-iodide, bromate ions, and water. This is the reaction in which we can determine the amount of I2 produced over a specific amount of time. The other two reactions are the reactions where the di-iodide will be consumed, and this is where the starch indicator produces a deep blue or violet color to indicate that a starch is present in the solution and the reaction has come to completion. As stated in the lab manual, every three moles of I2 produced consume six moles of thiosulfate (S2O32-), and the equation is given:

The second part of the experiment concentrates on the amount of energy required for a reaction to occur, depicted as Ea, and also known as the activation energy. As the temperature of the surroundings increase, more heat or energy is applied to the system, and in turn the reactions occur more quickly. This knowledge can be used to state that the rate of a reaction is dependent upon the temperature of the system and its surroundings.
Having determined the rate constant, k, the Arrhenius equation can be used to determine the activation energy necessary at a given temperature. By taking the natural log of both sides of the equation and plotting ln(k) and (1/T), the slope will equal (-Ea/R), which is the activation energy necessary in Joules or kilojoules required for the reaction to occur.