Chemical Equilibrium in Solution

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Chemical Equilibrium in Solution

2/5/13 Jim Jacobs (Partner: Jessica Janiga)

The purpose of this experiment was to determine the equilibrium established between iodine species [I3-, I2, and I-] dissolved in an aqueous KI solution. The values obtained for equilibrium were as follows: Kc (run#1) = 620; Kc (run#2) = 599; Kc (run#3) = 545. In effort to distinguish between the species of iodine the distribution coefficient ratio between two immiscible solvents (dichloromethane and water) was determined at multiple concentrations.

I. Introduction

The purpose of this experiment is to determine the equilibrium constant of iodine species at various concentrations. The equilibrium equation of the iodine species is as follows:
I2+I_=I3- (1) The equilibrium constant is calculated as:

KC=(I3-)I2I- (2)

Where the equilibrium constant KC is calculated from the concentrations of the species and does not take into consideration the activities of the species. Knowing Kc will indicate if the reaction favors products (if Kc>1) or if it favors products (Kc<1). Sodium thiosulfate is used to determine the concentration of Iodine in solution via titration in the following reaction:

2S2O32-+ I3- → S4O62-+ 3I-

Before titration excess KI is added to each solution to ensure that I2 is converted completely to I3-. This is done to ensure that molecular iodine is not loss to evaporation while in the aqueous phase (I3- is non-volatile). However, this method only exposes the total amount of iodine dissolved in the aqueous solution and not the concentrations of individual species. That determination can be aided by examination of the heterogeneous equilibrium existing between the aqueous layer and an organic layer comprised of dichloromethane. The distribution of molecular iodine between the two immiscible layers can be related by a…...

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