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Method of Initial Rates Iodine Clock

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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.

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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 solution, and about 20 mL of 0.2 percent soluble starch solution was placed into a 500mL volumetric flask, then distilled water was added until the solution reached 500mL. The preparation of Buffer B is as follows: 50 mL of 0.75 M NaAc solution, 100 mL of 0.22 M HAc solution, and about 10 mL of 0.2 percent soluble starch solution was placed into a 250-mL volumetric flask, then distilled water was added until the solution reach 250mL. In this experiment, four different 100 mL solutions were prepared and the reaction rates were measured for each. For Solution I: 5.0 mL of H3AsO3 was placed in a 500-mL beaker using a pipette along with 5.0 mL of KIO3, 65.0 mL of buffer A, and 25.0 mL of KI. For Solution II: 5.0 mL of H3AsO3 was placed in a 500-mL beaker using a pipette along with 10.0 mL of KIO3, 60.0 mL of buffer A, and 25.0 mL of KI. For Solution III: 5.0 mL of H3AsO3 was placed in a 500-mL beaker using a pipette along with 5.0 mL of KIO3, 40.0 mL of buffer A, and 50.0 mL of KI. For Solution IV: 5.0 mL of H3AsO3 was placed in a 500-mL beaker using a pipette along with 5.0 mL of KIO3, 65.0 mL of buffer B, and 25.0 mL of KI.

Data
Reaction MixtureH3AsO3IO3-Buffer ABuffer BI-Elapsed
Reaction TimeRate of Reaction
I5.0 mL5.0 mL65.0 mL25.0 mLN/A11.5*
II5.0 mL10.0 mL60.0 mL25.0 mL186013.2*
III5.0 mL5.0 mL40.0 mL50.0 mL11818.7*
IV5.0 mL5.0 mL65.0 mL25.0 mL146594.2*

Reaction MixtureInitial concentration of IInitital Concentration of IO3 I.2.1
II.2.05
III.1.1
IV.2.1

Discussion
Out of all of the proposed mechanisms for the reaction, our results most closely resemble the one proposed by Bray and Liebhafsky. The key factor in
this determination is that our color change in the experiment occurred very slowly and did not change to blue, but instead to yellow in all of our solutions. Bray and Liebhafsky’s mechanism is fast to equilibrate to an iodidic acid intermediate which is then slow to react to hypoiodous acid and dihypoiodous acid. The availability of acidic intermediates, combined with a most likely more acidic than intended solvent environment could stall the reaction, in addition to change the interaction of the product with the indicator due to hydrogen bonding interactions with the highly acidic environment. This would explain why our solution was yellow instead of blue and why our actual data took about five times longer to collect than expected. To correct these mistakes, we could be much more careful when preparing solutions in order to ensure proper pH. However, using data from a previous group we were able to calculate that the overall reaction is second order in nature.

Cite this Method of Initial Rates Iodine Clock

Method of Initial Rates Iodine Clock. (2016, Oct 08). Retrieved from https://graduateway.com/method-of-initial-rates-iodine-clock/

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