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).
Procedure
1. Using 100 mL graduated cylinders, mix 50.0 mL of 0.100 M Fe(NO3)2 and 50.0 mL of 0.100 M
AgNO3 in a flask. Stopper, shake well, and let the mixture stand for about 30 minutes, with
occasional shaking (once every 5 minutes or so).
2. While waiting for the reaction mixture to reach equilibrium, prepare a burette with 0.100 M KSCN.
3. When the reaction between Fe(NO3)2 and AgNO3 has reached equilibrium, pipette a 25.0 mL sample
of the supernatent liquid (liquid above the precipitate) into a 250 mL flask, taking care not to stir
up the solid at the bottom of the reaction mixture.
4. Titrate the 25.0 mL sample with KSCN solution. (It will probably take less than 10 mL, so don't
overshoot the endpoint.) The titration reaction is:
Ag+ + SCN–  AgSCN(s)
... (2)
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 Fe3+ produced in the equilibrium reaction to form the complex ion
FeSCN2+, which has an intense red colour:
Fe3+ + SCN–  FeSCN2+
... (3)
Therefore, the Fe3+ 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 FeSCN2+ were present, but at low
concentrations the coloration is orange.)
5. Repeat the titration with a second 25.0 mL portion of the equilibrium mixture. If the results do not
agree to within ± 0.05 mL, check with your instructor before repeating the titration a third time.
Determine the average volume of the two volumes of KSCN used which agree most closely with
each other. If you have three values that differ from each other by more or less equal amounts,
average all three values.
Table 1: Titration Data
Trial 1
Trial 2
Trial 3
Trial 4
Initial KSCN Burette Reading (mL)
Final KSCN Burette Reading (mL)
Volume of KSCN used (mL)
Calculations
1. Recall that when two solutions mix, they dilute each other as follows.
[solution #1]DILUTED = [solution #1]BEFORE MIXING x
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).