Redox Titrations

Specification Reference 5.2.3 (d)–(e)

Quick Notes

Redox titrations use oxidation–reduction reactions to determine unknown concentrations.
Common examples:
- Fe²⁺ with MnO₄⁻ in acidic conditions: Purple MnO₄⁻ is reduced to colourless Mn²⁺.
- I₂ with S₂O₃²⁻: Iodine (brown) is reduced to I⁻ (colourless).
Indicators are often unnecessary if the titrant is self-indicating (e.g. MnO₄⁻).
Balanced half-equations are used to find overall equations and mole ratios.
Calculations follow standard titration methods using volume, concentration, and mole ratios.

Full Notes

What Are Redox Titrations?

Redox titrations involve a redox reaction between a titrant and an analyte to determine concentration. They work in the same way as acid–base titrations (see titrations) however these reactions are used when the species involved undergo a change in oxidation state.

The procedure follows standard titration steps:

A solution of known concentration is added slowly from a burette.
The analyte reacts with the titrant in a redox reaction.
A colour change signals the endpoint.

Common Examples

1. Iron(II) and Manganate(VII) – Fe²⁺ / MnO₄⁻

Titrant: Potassium manganate(VII) (KMnO₄)
No indicator is needed – MnO₄⁻ is purple and turns colourless when reduced.
Half-equations:
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
Fe²⁺ → Fe³⁺ + e⁻
Overall reaction:
MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O
Usually performed in acidic solution (sulfuric acid is preferred to avoid side reactions with HCl or HNO₃).

2. Iodine and Thiosulfate – I₂ / S₂O₃²⁻

Iodine is generated in situ, often by oxidising iodide ions.
Titrant: Sodium thiosulfate (Na₂S₂O₃)
Half-equations:
I₂ + 2e⁻ → 2I⁻
2S₂O₃²⁻ → S₄O₆²⁻ + 2e⁻
Overall reaction:
I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻
To improve endpoint detection: starch indicator is used near the endpoint (deep blue with iodine, disappears when all I₂ is reduced).

Redox Titration Calculations

Write the balanced overall redox equation.
Calculate moles of titrant added using: n = c × V (in dm³).
Use mole ratios to find moles of analyte.
Calculate concentration or mass if required.

Calculations in Redox Titrations

You may be asked to:

Use moles, volumes, and concentrations to find unknown quantities.
Apply stoichiometry from the redox equations.

Worked Example

Worked Example – Finding the % of Iron in an Iron Tablet

Problem:
An iron tablet was dissolved and made up to 250.0 cm³. 25.0 cm³ of this solution was titrated with 0.0200 mol dm⁻³ KMnO₄. The average titre was 23.60 cm³. The tablet’s mass was 2.50 g (larger than before). Calculate the percentage by mass of iron (Fe²⁺) in the tablet. (Relative atomic mass of Fe = 55.8)

Step 1: Write the redox equation
MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O
Step 2: Calculate moles of MnO₄⁻ used
Moles = concentration × volume (in dm³)
Moles MnO₄⁻ = 0.0200 × (23.60 ÷ 1000) = 4.72 × 10⁻⁴ mol
Step 3: Find moles of Fe²⁺
From the stoichiometry, 1 mol MnO₄⁻ reacts with 5 mol Fe²⁺.
Thus: Moles Fe²⁺ = 5 × 4.72 × 10⁻⁴ = 2.36 × 10⁻³ mol (in 25.0 cm³)
Step 4: Scale up to 250.0 cm³
The whole solution is 10 times larger than the sample.
Total moles Fe²⁺ = 2.36 × 10⁻³ × 10 = 2.36 × 10⁻² mol
Step 5: Calculate mass of Fe
mass = moles × M_r
mass of Fe = 2.36 × 10⁻² × 55.8 = 1.317 g
Step 6: Find % of iron in the tablet
% Fe = (mass of Fe ÷ mass of tablet) × 100
% Fe = (1.317 ÷ 2.50) × 100 = 52.7%

Matt’s exam tip

Look out for dilutions with redox style titration questions. In this example, we have to remember the whole solution is 10× larger than the sample used in the titration. This is very common in these kind of exam questions.

Summary

Redox titrations use redox reactions to find unknown concentrations.
KMnO₄ is self-indicating and turns colourless when reduced to Mn²⁺.
Iodine–thiosulfate titrations use starch near the endpoint for a sharp colour change.
Balance half-equations to get overall equations and correct mole ratios.
Use n = c × V and stoichiometry to calculate unknowns.