Chemistry/Equilibrium Titration
The object of this experiment is to determine the value of the equilibrium constant for the following reaction using a titration procedure.
(1)
We will start with a solution containing equal concentrations of Ag^{+} and Fe^{2+} 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 Fe^{2+} which react. Also, the moles of Fe^{3+} which form equal the moles of Fe^{2+} 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).
Procedure
That is, the number of moles of Ag^{+} present at equilibrium equals the number of moles of SCN^{–} which are added from the burette. When all of the Ag^{+} has been removed from solution, the addition of one excess drop of SCN^{–} causes the Fe^{3+} produced in the equilibrium reaction to form the complex ion FeSCN^{2+}, which has an intense red colour:
Therefore, the Fe^{3+} acts as an indicator: when the solution being titrated takes on an orange tint, the titration is at the endpoint. (The solution would be red if sufficient FeSCN^{2+} were present, but at low concentrations the coloration is orange.)

Calculations
1. Recall that when two solutions mix, they dilute each other as follows.
[solution #1]DILUTED = [solution #1]BEFORE MIXING [math] * \frac {Volume\ of\ solution\ 1} {Volume\ of\ mixture} [/math]
When the AgNO_{3} solution is mixed with the Fe(NO_{3})_{2} solution, the two solutions dilute each other.
(a) Calculate the [Ag^{+} ] after the dilution.
(b) Calculate the [Fe^{2+} ] after the dilution.
These values represent the "starting [Ag^{+} ] and [Fe^{2+} ] 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 [Fe^{2+} ] DECREASE at the same time, according to equation (1)? Based on this "change in [Fe^{2+} ]" and the starting [Fe^{2+} ], found in calculation 1(b), what is the equilibrium [Fe^{2+} ]?
(c) What is the starting [Fe^{3+} ] (before the reaction starts)? NOTE: [Fe^{3+} ], NOT [Fe^{2+} ] ! According to equation (1), by how much does the [Fe^{3+} ] INCREASE when the reaction comes to equilibrium? What is the equilibrium [Fe^{3+} ]?
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).