Redox titrations

Why redox titrations suit this standard

• The mole ratios are almost never 1 : 1, so the stoichiometry does real work in the calculation.
• Several are self-indicating, so the endpoint chemistry can itself be explained and justified.
• The reactions are genuine Level 3 chemistry, with half equations that can be linked to the procedure.

The three common systems

Permanganate titrations

$$\text{MnO}_4^-(aq) + 8\text{H}^+(aq) + 5e^- \rightarrow \text{Mn}^{2+}(aq) + 4\text{H}_2\text{O}(l)$$

• Self-indicating — purple MnO₄⁻ becomes colourless Mn²⁺, so the endpoint is the first permanent faint pink.
• Must be acidified with sulfuric acid, never hydrochloric: chloride is a reductant and would be oxidised to chlorine in a competing reaction.
• Used for iron(II), oxalate, hydrogen peroxide.

Against iron(II), the overall equation gives the 1 : 5 ratio:

$$\text{MnO}_4^- + 8\text{H}^+ + 5\text{Fe}^{2+} \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} + 5\text{Fe}^{3+}$$

Iodine–thiosulfate titrations

$$\text{I}_2(aq) + 2\text{S}_2\text{O}_3^{2-}(aq) \rightarrow 2\text{I}^-(aq) + \text{S}_4\text{O}_6^{2-}(aq)$$

• Starch is the indicator, giving a sharp blue-black to colourless endpoint.
• Add the starch only near the endpoint, when the solution has faded to pale straw. Added early, starch binds iodine so strongly that the endpoint is sluggish and reads high.
• Used for ascorbic acid, hypochlorite in bleach, copper(II), and for iodine liberated from iodate.

Dichromate titrations

$$\text{Cr}_2\text{O}_7^{2-}(aq) + 14\text{H}^+(aq) + 6e^- \rightarrow 2\text{Cr}^{3+}(aq) + 7\text{H}_2\text{O}(l)$$

• Not self-indicating — the orange-to-green change is gradual, so a redox indicator is needed.
• Dichromate is stable enough to be used as a primary standard, which is its main advantage.
• It can be used with hydrochloric acid, unlike permanganate, because its potential is not high enough to oxidise chloride.

Choosing which to use

Analyte	Titrant	Ratio	Endpoint
Iron(II) in a supplement	acidified MnO4−	1 : 5	first permanent pink
Ascorbic acid (vitamin C)	I2 directly, or liberated I2 with S2O32−	1 : 1 with I2, then 1 : 2 with thiosulfate	blue-black → colourless
Hypochlorite in bleach	excess I−, then S2O32−	1 : 2	blue-black → colourless
Hydrogen peroxide	acidified MnO4−	2 : 5	first permanent pink
Oxalate	acidified MnO4−, warmed	2 : 5	first permanent pink

Justifying the procedure — the Excellence criterion

Excellence asks you to justify the steps used in the procedure in relation to the reactions occurring and to the nature of the samples. Every step in a redox titration has a chemical reason:

Step	The reaction-based justification
Acidify with sulfuric acid	the permanganate half equation consumes 8H+ per MnO4−; chloride from HCl would be oxidised in competition
Add a large excess of acid	8 mol of H+ per mole of permanganate is needed, and a shortfall gives brown MnO2 instead of colourless Mn2+
Warm an oxalate titration	high activation energy makes the reaction too slow at room temperature
Add starch late in an iodine titration	starch–iodine binding is strong, so early addition makes the endpoint sluggish and high
Add excess iodide in an iodometric determination	ensures all the oxidant is converted to an equivalent amount of I2
Titrate promptly after liberating iodine	iodine is volatile and is lost by evaporation, and iodide is air-oxidised
Stopper the flask between steps	limits both iodine loss and atmospheric oxidation

• Notice that each justification names a specific reaction or property, not a general principle. That is what "in relation to the reactions occurring" means.