Redox titrations using potassium manganate(VII) and iodine-thiosulfate, the calculations involved, colorimetry for coloured ions, and the planning and evaluation of quantitative analysis with attention to uncertainty.

What this dot point is asking

CCEA wants you to carry out and calculate redox titrations (manganate(VII) and iodine-thiosulfate), use colorimetry for coloured ions, and plan and evaluate quantitative analysis with attention to uncertainty.

Redox titrations in outline

A redox titration finds an unknown amount of an oxidising or reducing agent by reacting it with a standard solution of the other in a known mole ratio taken from the balanced ionic equation. The key skill is the mole calculation: find the moles of the standard from $n = c \times V$, scale by the ratio from the equation, then divide by the volume to get the unknown concentration. CCEA examines two systems in detail, manganate(VII) and iodine-thiosulfate, each with its own way of detecting the endpoint.

Manganate(VII) titrations

The manganate(VII) is placed in the burette and run into the acidified iron(II) solution. While any iron(II) remains, each drop of purple manganate(VII) is reduced to almost colourless $\text{Mn}^{2+}$ and the colour vanishes on swirling. Once all the iron(II) has reacted, the next drop is not reduced and the solution turns a permanent pale pink, which is the endpoint. The acid must be sulfuric: hydrochloric acid would be oxidised to chlorine by the manganate(VII), consuming extra oxidant and giving a falsely high titre, while nitric acid is itself an oxidising agent and would interfere. The amount of acid must be in excess so the medium stays strongly acidic, otherwise brown $\text{MnO}_2$ forms instead of the colourless $\text{Mn}^{2+}$.

Iodine-thiosulfate titrations

This is an example of an indirect (back) titration. The oxidising agent itself may have no convenient colour change, so it is first reacted with excess potassium iodide; the iodine set free is exactly equivalent to the oxidising agent, and the iodine is then measured by titration against thiosulfate. For example, copper(II) liberates iodine: $2\text{Cu}^{2+} + 4\text{I}^- \rightarrow 2\text{CuI} + \text{I}_2$, so each mole of $\text{Cu}^{2+}$ frees half a mole of $\text{I}_2$ and needs one mole of thiosulfate. As the brown iodine colour fades to pale straw the endpoint is close; only then is starch added, turning the mixture blue-black, and the endpoint is the single drop that makes the blue-black just disappear.

Colorimetry

The colorimeter passes light of a chosen colour (a filter is selected so the solution absorbs strongly, usually the complementary colour of the solution) through the sample, and the detector reads the absorbance. A series of standards of known concentration is measured to plot a calibration curve of absorbance against concentration, which is a straight line through the origin at low concentrations (the Beer-Lambert relationship). The absorbance of the unknown is then read against this line. Colorimetry suits coloured transition-metal ions and is faster and uses less sample than a titration.

Planning and uncertainty

Good quantitative analysis controls uncertainty. The percentage uncertainty in a measuring step is the absolute uncertainty divided by the reading, so using larger volumes and masses lowers the percentage uncertainty. A burette read to $\pm 0.05\ \text{cm}^3$ at each end gives $\pm 0.10\ \text{cm}^3$ on a titre, so a titre of $25\ \text{cm}^3$ carries a smaller percentage uncertainty than one of $10\ \text{cm}^3$. Repeat until titres are concordant (agreeing within $0.10\ \text{cm}^3$), and mean only the concordant results, discarding the rough trial. The overall percentage uncertainty of the experiment is found by adding the percentage uncertainties of each apparatus reading, and the largest contributor is usually the burette.

Examples in context

Manganate(VII) titrations measure the iron content of an ore or of an iron tablet: the tablet is dissolved in sulfuric acid and titrated, and the self-indicating endpoint avoids any added indicator. Iodine-thiosulfate titrations measure the chlorine content of pool water or the copper content of an alloy (copper(II) liberates iodine from iodide). CCEA practical assessment rewards quoting the percentage uncertainty in each measuring step and identifying the largest contributor, usually the burette reading.

Try this

Q1. State why manganate(VII) needs no separate indicator. [1 mark]

• Cue. It is self-indicating; the first permanent pink colour marks the endpoint.

Q2. Write the equation for the reaction of iodine with thiosulfate. [1 mark]

• Cue. $\text{I}_2 + 2\text{S}_2\text{O}_3^{2-} \rightarrow 2\text{I}^- + \text{S}_4\text{O}_6^{2-}$.

Q3. $20.0\ \text{cm}^3$ of $0.0100\ \text{mol dm}^{-3}$ manganate(VII) reacts with iron(II) in a $1:5$ ratio. Calculate the moles of iron(II). [2 marks]

• Cue. $n(\text{MnO}_4^-) = 2.00 \times 10^{-4}$; $n(\text{Fe}^{2+}) = 1.00 \times 10^{-3}\ \text{mol}$.