Redox Reactions and Qualitative Analysis

Study Notes

The reaction between Fe²⁺(aq) and I⁻(aq) results in Fe²⁺ being reduced to Fe²⁺ and I⁻ being oxidised to I₂.
The equation for this reaction is: 2Fe²⁺(aq) + 2I⁻(aq) → 2Fe²⁺(aq) + I₂(aq)
The reaction involves a colour change, from orange-brown to pale green.

Standard electrode potentials can be used to explain why redox reactions occur.
A more positive E value indicates that the equilibrium is more likely to gain electrons, shifting to the right and undergoing reduction.
Table 1 shows the standard electrode potentials for iron redox reactions.

The reaction between Fe²⁺ and MnO₄⁻ in acid conditions results in Fe²⁺ being oxidised to Fe³⁺ and MnO₄⁻ being reduced to Mn²⁺.
The E value for MnO₄⁻/Mn²⁺ is more positive than for Fe³⁺/Fe²⁺, so Fe²⁺ is oxidised to Fe³⁺ and MnO₄⁻ is reduced to Mn²⁺.
The reaction between Fe³⁺ and I⁻ results in I⁻ being oxidised to I₂ and Fe³⁺ being reduced to Fe²⁺.
The E value for Fe³⁺/Fe²⁺ is more positive than for I₂/I⁻, so I⁻ is oxidised to I₂ and Fe³⁺ is reduced to Fe²⁺.

Demonstrate knowledge and understanding of redox reactions and accompanying colour changes for the Fe²⁺/Fe³⁺ and Cr³⁺/Cr₂O₇²⁻ systems.
Apply knowledge of redox and disproportionation reactions of copper.
Perform qualitative analysis of ions on a test-tube scale.

The redox reaction between Fe²⁺ and MnO₄⁻ in acid conditions is used as a basis for a redox titration.
The equation for the reaction is: MnO₄⁻(aq) + 8H⁺(aq) + 5Fe²⁺(aq) → Mn²⁺(aq) + 5Fe³⁺(aq) + 4H₂O(l)
The solution containing MnO₄⁻ ions is purple and is decolourised by Fe²⁺(aq) ions to form a colourless solution containing Mn²⁺(aq) ions.