transition metals and redox titrations

what are redox titrations used for?

to find the concentration of an oxidising/reducing agent in solution (i.e. how much oxidising agent is needed to exactly react with a quantity of reducing agent)

why do you not need an additional indicator for a redox titration? give an example:

• redox titrations are self indicating

• purple potassium manganate (VII) is reduced to pale pink manganese (II) ions

give the ½ eqn for the reduction of potassium manganate (VII) → manganese (II) ions:

MnO₄^- + 8H⁺ + 5e^- → Mn²⁺ + 4H₂

give the method for a typical redox titration:

1. measure out the quantity of reducing agent using a pipette (either Fe²⁺ or C₂O₄^2- ions)

2. add dilute sulfuric acid so it is in excess

3. add the oxidising agent to the burette (usually potassium manganate)

4. add the oxidising agent from the burette until the solution just changes colour (to the colour of the oxidising agent)

5. repeat to get concordant results

6. calculate the concentration of the reducing agent

what is a reducing agent? give two examples and compounds in which they may be found:

e^- donor:

• Fe²⁺ ions - iron (II) sulfate, ammonium iron (II) sulfate

• C₂O₄^2- ions - sodium oxalate, potassium oxalate

give the ½ eqn for the oxidation of Fe²⁺ → Fe³⁺:

Fe²⁺ → Fe³⁺ + e^-

give the ½ eqn for the oxidation of C₂O₄^2- → CO₂:

C₂O₄^2- → 2CO₂ + 2e^-

give the eqn for the reaction of Fe²⁺ with MnO₄^-:

MnO₄^- _(aq) + 8H⁺ _(aq) + 5Fe²⁺ _(aq) → Mn²⁺ _(aq) + 4H₂O _(l) + 5Fe³⁺ _(aq)

give the eqn for the reaction of C₂O₄^2- with MnO₄^-:

2MnO₄^- _(aq) + 16H⁺ _(aq) + 5C₂O₄^2- _(aq) → 2Mn²⁺ _(aq) + 8H₂O _(l) + 10CO_{2 (g)}