Writing Half Equations for Electrode Reactions Revision Notes | GCSE AQA Higher Chemistry

1. What Are Half Equations?

During electrolysis, different reactions happen at each electrode. At one electrode, substances gain electrons. At the other, substances lose electrons. Half equations let us write these reactions separately, showing exactly what happens to the electrons.

Think of half equations as snapshots of each electrode's story. The full electrolysis is like a dance - electrons leave one partner (oxidation) and join another (reduction). Half equations capture each side of this electron transfer.

2. The Two Types of Half Equations

There are only two types of half equations you need to master. At the cathode, substances gain electrons (reduction). At the anode, substances lose electrons (oxidation). The memory trick: AN OX, RED CAT - ANode = OXidation, REDuction = CAThode.

Let's see this with the example from the specification. When hydrogen ions gain electrons at the cathode: $2H^+ + 2e^- \rightarrow H_2$. This is reduction because electrons are gained. When hydroxide ions lose electrons at the anode: $4OH^- \rightarrow O_2 + 2H_2O + 4e^-$. This is oxidation because electrons are lost.

3. Writing Half Equations Step by Step

Writing half equations follows a clear method. You need to balance atoms first, then balance charge by adding electrons. This catches many students out because they forget the charge balance step.

Worked example: Writing a Half Equation for Chloride Ion Oxidation

Write the half equation for chloride ions forming chlorine gas at the anode

1. $Cl^- \rightarrow Cl_2$ — Start with reactant and product, ignore electrons for now
2. $2Cl^- \rightarrow Cl_2$ — Need 2 chloride ions to make 1 chlorine molecule
3. Left side: $2 \times (-1) = -2$ Right side: $0$ — Charges are not balanced - left side is more negative
4. $2Cl^- \rightarrow Cl_2 + 2e^-$ — Add 2 electrons to right side to balance the -2 charge on left

Answer: $2Cl^- \rightarrow Cl_2 + 2e^-$

Notice how electrons appear on the product side? That tells us this is oxidation (loss of electrons). The electrons have to go somewhere - they travel through the external circuit to the cathode where reduction happens.

4. Common Half Equations You Must Know

The specification gives you key examples that appear in exams. These are the building blocks for more complex electrolysis questions. Let's break down each one and understand when they occur.

Both oxygen equations show the same reaction - it is a different way of writing where the electrons go. The first shows electrons as products (gained by the external circuit). The second shows electrons being subtracted (lost by the hydroxide ions). Examiners accept both formats.

5. Balancing Charge - The Critical Step

This is where most students lose marks. You must balance both atoms AND charge. Many students balance atoms correctly but forget about charge balance. The electrons are what make the charges balance.

Worked example: Balancing a Tricky Half Equation

Write the half equation for aluminium ions forming aluminium metal at the cathode

1. $Al^{3+} \rightarrow Al$ — Aluminium ions become aluminium metal
2. 1 Al atom on each side ✓ — Atoms are already balanced
3. Left: $+3$, Right: $0$ — Left side is +3 more positive than right side
4. $Al^{3+} + 3e^- \rightarrow Al$ — Add 3 electrons to left side to neutralise the +3 charge
5. Atoms: 1 Al = 1 Al ✓ Charge: $(+3) + 3(-1) = 0$ ✓ — Both atoms and charge are balanced

Answer: $Al^{3+} + 3e^- \rightarrow Al$

The key insight: electrons carry negative charge. If your left side is more positive, add electrons to the left. If your right side is more positive, add electrons to the right. The electrons always go where they're needed to balance the charge.

6. Exam Technique and Common Mistakes

Examiners have favourite tricks when testing half equations. They love to test whether you can spot incorrectly balanced equations, and they often ask you to identify oxidation versus reduction. Here's how to get full marks every time.

The biggest trap: students write $H^+ + e^- \rightarrow H_2$ instead of $2H^+ + 2e^- \rightarrow H_2$. One hydrogen ion cannot make a hydrogen molecule - you need two. Always check your atom balance first, then your charge balance.

Another common error: confusing which electrode is which. The cathode is always where reduction happens (electrons gained). The anode is always where oxidation happens (electrons lost). This is true for all electrolysis, whether the electrolyte is molten or aqueous.

Half-equation exam tip: In Writing Half Equations for Electrode Reactions - Representation of reactions at electrodes as half equations (HT only) - Electrolysis, always identify the key chemical idea first, then link it to evidence, data, or an equation. A high-scoring GCSE answer names the particles involved, states what changes, and justifies the conclusion using correct scientific language. Where calculations are used, show each step with units and keep significant figures sensible. Finally, check that terminology matches AQA wording so your answer is unambiguous and mark-scheme friendly.

A half equation must balance both atoms and charge. That is why electrons are written explicitly: they show the transfer taking place at the electrode. Reduction half equations gain electrons, while oxidation half equations show electrons being released.