Oxidation and reduction defined in terms of electron loss and gain, the meaning of oxidising and reducing agents, writing ion-electron half-equations, combining them into redox equations, and the electrochemical series.

What this key area is asking

The SQA wants you to define oxidation and reduction in terms of electrons, identify oxidising and reducing agents, write and balance ion-electron half-equations (including those involving oxygen and hydrogen), combine them into overall redox equations, and use the electrochemical series to predict which way electrons flow. This is a calculation-rich key area: the redox titration appears almost every year, so the half-equation skills here feed directly into the analysis questions in Area 3.

Oxidation and reduction

An oxidising agent removes electrons from another species, so it is itself reduced. A reducing agent gives electrons to another species, so it is itself oxidised. Strong oxidising agents include the halogens and acidified permanganate ($MnO_4^-$); strong reducing agents include the Group 1 metals and the iron(II) ion $Fe^{2+}$.

Ion-electron half-equations

A half-equation shows either the loss or the gain of electrons:

• Oxidation of magnesium: $Mg \rightarrow Mg^{2+} + 2e^-$
• Reduction of chlorine: $Cl_2 + 2e^- \rightarrow 2Cl^-$

For half-equations involving oxygen, balance the atoms using water and hydrogen ions, then balance the charge with electrons. Many of these are given in the SQA data booklet, but you must be able to scale and combine them.

Combining half-equations

For example, combining $Mg \rightarrow Mg^{2+} + 2e^-$ with $Cl_2 + 2e^- \rightarrow 2Cl^-$ gives $Mg + Cl_2 \rightarrow Mg^{2+} + 2Cl^-$.

Worked example: a permanganate redox titration

The electrochemical series

The electrochemical series in the data booklet lists reduction half-equations. A species higher in the list is a stronger reducing agent on its left-hand side; species lower down are stronger oxidising agents. This lets you predict which way electrons flow when two half-cells are connected: electrons flow from the half-cell higher in the series (the stronger reducing agent) to the one lower down.

Examples in context

In water treatment, chlorine acts as a powerful oxidising agent: $Cl_2 + 2e^- \rightarrow 2Cl^-$ describes its reduction as it removes electrons from, and so destroys, the molecules in bacteria. The same electron-accepting behaviour explains why acidified permanganate is used in the lab to determine the iron content of an iron tablet or a sample of water. A pharmaceutical quality-control chemist dissolves a crushed iron tablet in dilute sulfuric acid and titrates the $Fe^{2+}$ against standardised permanganate, using exactly the $1 : 5$ ratio above to back-calculate the iron(II) content per tablet. The self-indicating purple-to-colourless change is why no separate indicator is added, a detail SQA markers reward.

Try this

Q1. Write the ion-electron half-equation for the oxidation of iron(II) to iron(III). [1 mark]

• Cue. $Fe^{2+} \rightarrow Fe^{3+} + e^-$.

Q2. In a reaction, zinc displaces copper from copper(II) sulfate. Identify the reducing agent. [1 mark]

• Cue. Zinc, because it loses electrons (is oxidised) and donates them to the copper ions.

Q3. $20.0 \text{ cm}^3$ of iron(II) solution needs $16.0 \text{ cm}^3$ of $0.0200 \text{ mol l}^{-1}$ permanganate. Calculate the iron(II) concentration. [3 marks]

• Cue. $n(MnO_4^-) = 3.20 \times 10^{-4}$ mol; $\times 5 = 1.60 \times 10^{-3}$ mol $Fe^{2+}$; $C = 1.60 \times 10^{-3} / 0.0200 = 0.0800 \text{ mol l}^{-1}$.