What is a redox titration and what is it used for?
Titration is a laboratory method that is used to determine the concentration or mass of a substance (called the analyte). A solution of known concentration, called the titrant, is added to a solution of the analyte until just enough has been added to react with all of the analyte (the equivalence point). If the reaction between the titrant and the analyte is a reduction-oxidation reaction, the procedure is called a redox titration.
One example is the use of potassium permanganate to determine the percentage of iron in an unknown iron(II) salt.
The equation for the reaction is
MnO₄⁻ + 5Fe²⁺ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O
The Fe³⁺ ion has a pale yellow-green colour, and the Mn²⁺ has a pale orange-pink colour, but the deep purple colour of MnO₄⁻ is so intense that it is the colour you see at the equivalence point.
In the titration, you place a known mass (say, 1.2352 g) of the sample, 10 mL of water, and 10 mL of 1 mol/L H₂SO₄ in a conical flask. Then you slowly add a solution of KMnO₄ of known concentration (say, 0.020 48 mol/L) from a buret. The MnO₄⁻ is decolourized immediately when it is first added, but the colour takes longer to fade as you approach the equivalence point. You have reached the equivalence point when one drop of MnO₄⁻ produces a pale permanent purple-pink colour (say, after adding 26.01 mL).
You know the volume and the molarity of the KMnO₄ and the mass of the unknown compound, so you can calculate the mass of the Fe²⁺.
Here’s how you do it.
Moles of MnO₄⁻ = 0.026 01 L MnO₄⁻ × (0.020 48 mol MnO₄⁻/1 L MnO₄⁻) =
5.327 × 10⁻⁴ mol MnO₄⁻
Moles of Fe²⁺ = 5.327 × 10⁻⁴ mol MnO₄⁻ × (5 mol Fe²⁺/1 mol MnO₄⁻) =
2.663 × 10⁻³ mol Fe²⁺
Mass of Fe²⁺ = 2.663 × 10⁻³ mol Fe²⁺ × (55.845 g Fe²⁺/1 mol Fe²⁺) = 0.1487 g Fe²⁺
% Fe²⁺ = 0.1487 g Fe²⁺/1.2352 g sample × 100 % = 12.04 %
You use a titration to determine the concentration or amount of a substance in an unknown sample.