Redox titrationsEdexcel International A Level Chemistry: Revision notes

In a redox titration a solution of an oxidising agent reacts with a reducing agent (or the reverse) and the volume needed to reach the end-point is used to find an unknown concentration or formula.

The reaction must be fast, go to completion and have a clear end-point. Calculations use the balanced equation: find moles of the known solution (c × V/1000), apply the mole ratio, then work out the unknown. Titres are repeated until concordant (within 0.10 cm³) and the mean of concordant titres is used.

Acidified potassium manganate(VII) oxidises iron(II) ions:

MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O

MnO₄⁻ is reduced (Mn +7 to +2) and Fe²⁺ is oxidised (+2 to +3). The mole ratio is 1 : 5.

• Manganate(VII) is in the burette and iron(II) in the conical flask with an excess of dilute sulfuric acid.
• It is its own indicator: the end-point is the first permanent pale pink.
• Do not use hydrochloric acid, because MnO₄⁻ would oxidise Cl⁻ to Cl₂ and the titre would be too large. Do not use nitric acid, which is itself an oxidising agent.
• Too little acid gives brown MnO₂.

25.0 cm³ of Fe²⁺ solution needs 22.40 cm³ of 0.0200 mol dm⁻³ MnO₄⁻.

1. Moles MnO₄⁻ = 0.0200 × 22.40/1000 = 4.48 × 10⁻⁴ mol
2. Moles Fe²⁺ = 5 × 4.48 × 10⁻⁴ = 2.24 × 10⁻³ mol
3. Concentration of Fe²⁺ = 2.24 × 10⁻³ / 0.0250 = 0.0896 mol dm⁻³

If the 25.0 cm³ was a sample from 250 cm³ made up in a volumetric flask, multiply the moles by 10 to get the total in the flask. Then, for hydrated crystals FeSO₄·xH₂O, find M = mass / moles and x = (M − 151.9) / 18.0.

Sodium thiosulfate reduces iodine:

I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻

The mole ratio is 1 : 2. Iodine is reduced (0 to −1) and sulfur in thiosulfate is oxidised (+2 to +2.5).

• Thiosulfate is in the burette; the iodine solution is in the flask and is brown, then fades to pale yellow.
• Add starch only near the end-point (when pale yellow). The solution turns blue-black and the end-point is the first colourless solution.
• Adding starch too early makes the starch–iodine complex hold iodine and the end-point is inaccurate.

Example: 25.0 cm³ I₂ needs 18.60 cm³ of 0.100 mol dm⁻³ thiosulfate. Moles S₂O₃²⁻ = 1.86 × 10⁻³; moles I₂ = 9.30 × 10⁻⁴; [I₂] = 0.0372 mol dm⁻³.

Every piece of apparatus has an uncertainty. The percentage uncertainty is (absolute uncertainty / measured value) × 100.

• Burette: ±0.05 cm³ per reading, so ±0.10 cm³ for a titre (two readings).
• Pipette: e.g. ±0.06 cm³ on 25.0 cm³.
• Balance: ±0.005 g per reading, so ±0.01 g for a mass found by difference.

Add the percentage uncertainties for the total. Reduce the largest one: use a larger titre or larger mass.

For validity: use concordant titres, titrate Fe²⁺ solutions promptly (air oxidises Fe²⁺), use enough acid, and compare the result with the expected value within the uncertainty.

13a: redox titration of iron(II) with manganate(VII). Pipette 25.0 cm³ of Fe²⁺ into a conical flask, add excess dilute sulfuric acid, titrate with MnO₄⁻ to the first permanent pale pink, do a rough titre, then repeat to concordance.

13b: titration of iodine with thiosulfate, adding starch near the end-point and stopping at blue-black to colourless.

Good practice: rinse the burette with the solution it will hold, remove the funnel, read the bottom of the meniscus for thiosulfate (colourless) and the top for manganate(VII) (opaque purple), and swirl continuously.