Discussing the activity
The concentration of Ag+(aq) at equilibrium is found by titration with potassium thiocyanate solution. What is the ionic equation for this reaction?
Ag+(aq) + SCN–(aq) AgSCN(s)
Why don't we have to add an indicator to show the end point of the titration?
The equilibrium mixture contains Fe3+(aq) ions. When all of the Ag+(aq) ions have been preciptated as AgSCN(s), the SCN–(aq) ions react with Fe3+(aq) ions to produce a red colour due to Fe(SCN)2+. The appearance of the red colour marks the end point of the titration of Ag+(aq) with SCN–(aq).
Why don't we have to carry out an experiment to find the equilibrium concentration of Fe2+(aq) ions?
Because the initial concentrations of Fe2+(aq) and Ag+(aq) are equal and the ions react together in a one to one ratio the equilibrium concentration of Fe2+(aq) must be the same as the equilibrium concentration of Ag+(aq), which we have found by experiment.
How do we work out the equilibrium concentration of Fe3+(aq) ions?
We know the initial concentration of Ag+(aq) ions and we have found the concentration of this ion at equilibrium by experiment. By subtraction we can work out how many moles of Ag+(aq) have been used up in achieving equilibrium. The same number of moles of Fe3+(aq) must have also been produced when equilibrium is established. This means we can work out the equilibrium concentration of Fe3+(aq).
What about the equilibrium concentration of Ag(s). Don't we need to work this out?
No. This is a heterogeneous equilibrium so we don't include the concentration of a pure solid in the expression for Kc.
How can you calculate Kc for this equilibrium?
Put the equilibrium concentrations of Fe2+, Fe3+ and Ag+ into the expression: