Permanganate Titration of Oxidised Iron(II) Sulfate
An impure sample of iron(II) sulfate heptahydrate, FeSO4·7H2O, has been stored in a loosely stoppered bottle for two days. A student weighs 1.20 g of the sample, dissolves it in dilute sulfuric acid, and makes the solution up to 250.0 cm3 in a volumetric flask. They titrate 25.0 cm3 portions with 0.00200 mol dm-3 potassium manganate(VII) solution. The mean titre is 18.20 cm3. The student assumes that all the iron in the sample is present as Fe2+ and calculates the percentage purity from the titration result. The teacher notes that some of the pale green crystals have developed a brown coating and that the titration endpoint is a faint permanent pink. The reaction in acid is: MnO4- + 8H+ + 5Fe2+ → Mn2+ + 5Fe3+ + 4H2O.
1. Calculate the apparent percentage purity of the sample, assuming all iron is present as Fe2+.
Read the worked answer
Moles of MnO4- used = 0.00200 × 0.01820 = 3.64 × 10^-5 mol. From the equation, 1 mol MnO4- reacts with 5 mol Fe2+, so moles Fe2+ in 25.0 cm3 = 5 × 3.64 × 10^-5 = 1.82 × 10^-4 mol. The 250.0 cm3 flask contains 10 times this amount: 1.82 × 10^-3 mol Fe2+. Assuming this all comes from FeSO4·7H2O, mass = 1.82 × 10^-3 × 278.0 = 0.506 g. Percentage purity = 0.506 / 1.20 × 100 = 42.2%. This is an apparent value because the assumption that all iron is Fe2+ may be false.
2. Explain how air oxidation affects the calculated purity. Include an equation for the oxidation of Fe2+ by oxygen in acidic solution.
Read the worked answer
Air oxidises some Fe2+ to Fe3+: 4Fe2+ + O2 + 4H+ → 4Fe3+ + 2H2O. The MnO4- titration detects only Fe2+, because Fe3+ is already in its higher oxidation state and is not oxidised further under these conditions. Oxidation therefore lowers the number of moles of Fe2+ titrated, so the calculated mass of FeSO4·7H2O and the apparent percentage purity are lower than the true values. The brown coating is consistent with iron(III) products. The result is an underestimate, not an overestimate.
3. Suggest a modification that would allow the total iron content to be determined, and explain why this gives a better measure of the original FeSO4·7H2O content.
Read the worked answer
Before titration, add zinc and acid to reduce any Fe3+ back to Fe2+, then titrate the total iron with MnO4-. The zinc reduces Fe3+ to Fe2+ without changing the total amount of iron. Since all the iron originally came from FeSO4·7H2O, measuring total iron gives the original iron(II) content even if some air oxidation has occurred. Alternatively, use a freshly opened sample, dissolve it in deoxygenated dilute acid, and titrate immediately to minimise oxidation. The reduction method is preferable for an already oxidised sample because it corrects for the Fe3+ present.