Chemistry/Equilibrium Titration

CHEMISTRY 12 DETERMINATION OF  THE  EQUILIBRIUM  CONSTANT

The object of this experiment is to determine the value of the equilibrium constant for the following reaction using a titration procedure.

Ag+ (aq) + Fe2+ (aq) <===> Fe3+ (aq) + Ag(s)

(1)

We will start with a solution containing equal concentrations of Ag+ and Fe2+ and allow the mixture to come to equilibrium according to equation (1). As can be seen from the reaction equation, the moles of Ag+ which react equals the moles of Fe2+ which react. Also, the moles of Fe3+ which form equal the moles of Fe2+ which react.

When equilibrium has occurred, a sample of the equilibrium mixture is reacted with SCN– solution until all the Ag+ remaining at equilibrium is precipitated as AgSCN(s).

Calculations

1. 	Recall that when two solutions mix, they dilute each other as follows.

[solution #1]DILUTED = [solution #1]BEFORE MIXING $$ * \frac {Volume\ of\ solution\ 1} {Volume\ of\ mixture} $$

When the AgNO3 solution is mixed with the Fe(NO3)2 solution, the two solutions dilute each other.

(a) 	Calculate the [Ag+ ] after the dilution.

(b) 	Calculate the [Fe2+ ] after the dilution.

These values represent the "starting [Ag+ ] and [Fe2+ ] in the solution".

2. 	Use the [KSCN] (that is, [SCN– ]) and the average volume of KSCN added in the titrations to calculate the moles of SCN– added.

3. 	(a) 	Based on equation (2), how many moles of Ag+ are present in the 25.0 mL sample of the equilibrium mixture you titrated? Based on the number of moles of Ag+ present in 25.0 mL of the mixture, what [Ag+ ] exists in the 25.0 mL sample at equilibrium?

(b) 	Using the starting [Ag+ ], found in calculation 1(a), and the [Ag+ ] existing at equilibrium, as found in calculation 3(a), calculate the decrease in [Ag+ ] occurring when the starting mixture reaches equilibrium. Note that we calculate changes in concentration in the same way that we calculate changes in moles: since all the concentrations refer to the same volume, then


 * equilibrium moles = starting moles – change in moles
 * and   	equilibrium [ ]   = starting [ ]   – change in [ ].

Now that you have calculated the "change in [Ag+ ]", by how much will the [Fe2+ ] DECREASE at the same time, according to equation (1)? Based on this "change in [Fe2+ ]" and the starting [Fe2+ ], found in calculation 1(b), what is the equilibrium [Fe2+ ]?

(c) 	What is the starting [Fe3+ ] (before the reaction starts)? NOTE: [Fe3+ ], NOT [Fe2+ ] ! According to equation (1), by how much does the [Fe3+ ] INCREASE when the reaction comes to equilibrium? What is the equilibrium [Fe3+ ]?

4. 	Write out the equilibrium expression for reaction (1).

5. 	Using the results you arrived at in Calculation 3, calculate the value of the equilibrium constant for equation (1).