Method of Initial Rates Iodine Clock
Lab #3: Method of Initial Rates: Iodine Clock
Introduction The detailed system of steps in a reaction is called the reaction mechanism, and it is one of the principal aims of chemical kinetics to obtain information to aid in the elucidation of these mechanisms in order to better understand chemical processes. Reactions usually occur in a stepwise manner with each step proceeding at a different speed. If the rate of reaction is slow enough to measure, this is indicative of a step much slower than the rest of the process, known as the rate limiting step. For most reactions, a steady reaction state is quickly attained in which the concentration of reaction intermediates becomes dependent on the rate limiting step and closely associated steps. From this principle, the common rate law equation was experimentally determined (Eq. 1). Experimentally, the rate of reaction can be determined by altering the concentrations of the reactants. For our experiment we used the method of initial rates.
The method of initial rates requires that several rates are determined along with several different combinations of the concentrations. For this reaction, the rate was determined by an abrupt color change of the solution from clear to blue. This sudden emergence of a blue color is evidence that the arsenious acid has been consumed and free iodine atoms are liberated.
The purpose of the experiment was to determine the rate constant K and the orders of each reactant. This experiment can be used as a model for determining reaction rates in a variety of solutions.
Materials
• 20 mL of 0.2% of soluble starch
• 250-mL volumetric flask
• 500-mL volumetric flask
• 0.03 M H3AsO3
• 0.1 M KIO3
• 0.2 M KI
Procedure Buffer A and buffer B were prepared first with the concentration of Buffer A being roughly half that of Buffer B. The preparation of Buffer A is as follows: 100 mL of 0.75 M NaAc solution, 100 mL of 0.22 M HAc
Introduction The detailed system of steps in a reaction is called the reaction mechanism, and it is one of the principal aims of chemical kinetics to obtain information to aid in the elucidation of these mechanisms in order to better understand chemical processes. Reactions usually occur in a stepwise manner with each step proceeding at a different speed. If the rate of reaction is slow enough to measure, this is indicative of a step much slower than the rest of the process, known as the rate limiting step. For most reactions, a steady reaction state is quickly attained in which the concentration of reaction intermediates becomes dependent on the rate limiting step and closely associated steps. From this principle, the common rate law equation was experimentally determined (Eq. 1). Experimentally, the rate of reaction can be determined by altering the concentrations of the reactants. For our experiment we used the method of initial rates.
The method of initial rates requires that several rates are determined along with several different combinations of the concentrations. For this reaction, the rate was determined by an abrupt color change of the solution from clear to blue. This sudden emergence of a blue color is evidence that the arsenious acid has been consumed and free iodine atoms are liberated.
The purpose of the experiment was to determine the rate constant K and the orders of each reactant. This experiment can be used as a model for determining reaction rates in a variety of solutions.
Materials
• 20 mL of 0.2% of soluble starch
• 250-mL volumetric flask
• 500-mL volumetric flask
• 0.03 M H3AsO3
• 0.1 M KIO3
• 0.2 M KI
Procedure Buffer A and buffer B were prepared first with the concentration of Buffer A being roughly half that of Buffer B. The preparation of Buffer A is as follows: 100 mL of 0.75 M NaAc solution, 100 mL of 0.22 M HAc