Iron(II) & Iron(III) Ions
Oxidation of Fe2+ to Fe3+
- A redox titration involves an oxidising agent being titrated against a reducing agent
- Electrons are transferred from one species to another
- In acid-base titrations indicators are used to show the endpoint of a reaction; however redox titrations using transition metal ions naturally change colour when changing oxidation state, so indicators are not always necessary
- They are said to be 'self-indicating'
- The redox reaction between iron(II) ions and manganate(VII) ions in acidic conditions is used as a basis for a redox titration
- In this reaction:
- Fe2+ is oxidised to Fe3+
- MnO4- is reduced to Mn2+
- Potassium manganate(VII) is commonly used which is an oxidising agent and is a deep purple colour
- In acidic solutions it is reduced to the almost colourless manganese(II) ion by the Fe2+(aq)
- The equation for the reaction is:
MnO4- (aq) + 8H+ (aq) + 5Fe2+ → Mn2+ (aq) + 5Fe3+ + 4H2O (aq)
purple colourless
Reduction of Fe3+ to Fe2+
- An orange-brown solution of Fe3+(aq) ions can be reduced to pale green Fe2+(aq) ions by various reducing agents.
- A potassium iodide solution is commonly used
- The colour change can be masked by the formation of iodine which has a brown colour
- In this reaction
- Fe3+ is reduced to Fe2+
- I- is oxidised to I2
- The equation for this reaction is:
2Fe3+ (aq) + 2I- (aq) → 2Fe2+ + I2 (aq)
orange-brown pale-green brown