CHEM 361 L

EXPERIMENT 1 :The Iodine Clock

(Last revision: January 25, 2005)

Introduction

This experiment demonstrates the determination of the kinetic parameters of a multi-step, condensed phase chemical reaction. The kinetic parameters will be determined using the method of initial rates. The overall reaction is:

IO₃^- + 8 I^- +6 H⁺ = 3 I₃^- + 3 H₂O

which will be followed by monitoring the appearance of free iodine (as I₃^-). Arsenious acid is used to react with the nascent iodine,

H₃AsO₃ + I₃- + H₂O = HAsO₄^2- + 3 I^- + 4 H⁺

and soluble starch is used as an indicator. Once the arsenious acid has been used up, the iodine reacts with the starch producing a blue colored complex. The time required for the appearance of the blue color will be measured as a function of pH and reagent concentration (constant temperature and ionic strength) to get the exponents in the rate equation.

Procedure

All groups will work at the same temperature (25 ^oC). Prepare two 500 mL aqueous buffer solutions (buffer one containing 100 mL 0.75 M NaAc and 100 mL 0.22 M HAc; buffer two containing 100 mL 0.75 M NaAc and 200 mL 0.22 M HAc), 250 mL 0.03 M H₃AsO₃ (made up from NaAsO₂, stabilized at pH 5 with HAc), 250 mL 0.1 M KIO₃, and 1.00 L 0.2 M KI. The solutions should be kept in 25 ^oC water baths until used. Each group will perform two runs with each of the four combinations of reagents listed below:

• run 1: 5 mL H₃AsO₃, 5 mL KIO₃, 65 mL buffer one, 25 mL KI solutions.
• run 2: 5 mL H₃AsO₃, 10 mL KIO₃, 60 mL buffer one, 25 mL KI solutions.
• run 3: 5 mL H₃AsO₃, 5 mL KIO₃, 40 mL buffer one, 50 mL KI solutions.
• run 4: 5 mL H₃AsO₃, 5 mL KIO₃, 65 mL buffer two, 25 mL KI solutions.

The reactions should be performed in 250 mL Erlenmeyer flasks in the water baths. All of the amounts should be added to the flask, except the KI solution. A measured amount of spray starch solution should be dissolved into the solution, followed by the addition of the KI solution. The timing begins when the KI is added. Stop timing when the blue color appears.

Calculations

Determine [H⁺] for each buffer from the stock solution concentrations, using the Debye- Huckel law to get the activity coefficients (the ionic strength for these solutions is 0.16). From the known initial concentration of arsenious acid, and the overall reaction stoichiometry, calculate the average rate of reaction (moles/L s) for each run. From combinations of runs, calculate the order of reaction (1) with respect to each of the reagents (round to integer or half- integer exponents). From the observed rates and orders, calculate an average rate constant based on all eight runs. Compare the results obtained with those in the literature, noting the degree of consistency with proposed mechanisms.