How to Study Electrochemistry for IGCSE

Electrochemistry is one of the most conceptually demanding topics in 0620. It combines ideas about ions, electron transfer, redox, and the reactivity series into a single framework, and students who have gaps in any of those foundation topics struggle here. The study strategy is therefore twofold: shore up the foundations first, then build the electrochemistry content on top.

Foundations you need before starting

Check that you can confidently do the following before studying electrochemistry in depth:

• Write the formula of any common ionic compound
• Identify cations and anions and their charges
• Explain what happens when an ionic compound melts or dissolves (ions become free to move)
• Place metals in reactivity series order
• Define oxidation and reduction in terms of electron transfer (OILRIG: Oxidation Is Loss, Reduction Is Gain)

If any of these are shaky, fix them first. Studying electrolysis without understanding ions is like studying equations without knowing element symbols.

Electrolysis: the three scenarios

The 0620 syllabus tests electrolysis in three main contexts. Study each one separately, then compare them.

1. Electrolysis of molten ionic compounds

This is the simplest case and the one to learn first.

• Only two ions are present (the cation and the anion from the compound).
• The cation moves to the cathode (negative electrode) and is reduced (gains electrons).
• The anion moves to the anode (positive electrode) and is oxidised (loses electrons).

Example: molten lead(II) bromide, PbBr2.

• At the cathode: Pb2+ + 2e- → Pb (lead metal deposited)
• At the anode: 2Br- → Br2 + 2e- (bromine gas released)

Practise writing these half-equations for molten NaCl, molten Al2O3, and molten MgCl2. The pattern is identical every time.

2. Electrolysis of aqueous solutions

This is harder because water introduces additional ions: H+ and OH-. Now there are competing ions at each electrode, and you need rules to predict which one reacts:

At the cathode: If the metal is above hydrogen in the reactivity series, hydrogen gas is produced (H+ ions are discharged). If the metal is below hydrogen, the metal is deposited.

At the anode: If the anion is a halide (Cl-, Br-, I-), the halogen is produced. If the anion is sulfate or nitrate, oxygen gas is produced (OH- ions are discharged instead).

These rules are the single most important piece of content in the topic. Card them on flashcards and test yourself until they are automatic.

Example: electrolysis of aqueous copper(II) sulfate with inert electrodes.

• Cathode: Cu2+ + 2e- → Cu (copper is below hydrogen, so copper is deposited)
• Anode: 4OH- → 2H2O + O2 + 4e- (sulfate is not a halide, so oxygen is produced)

Example: electrolysis of concentrated aqueous sodium chloride.

• Cathode: 2H+ + 2e- → H2 (sodium is above hydrogen, so hydrogen is produced)
• Anode: 2Cl- → Cl2 + 2e- (chloride is a halide, so chlorine is produced)

3. Electrolysis with non-inert electrodes

When the anode is made of a reactive metal (e.g., copper in copper sulfate solution), the anode itself dissolves rather than the anion being discharged. This is the basis of electroplating and copper purification. The anode loses mass, the cathode gains mass, and the solution concentration stays the same.

Study techniques for electrochemistry

Draw, do not just read

Electrochemistry is visual. For every electrolysis example, draw the cell:

• Label the battery, cathode (negative), anode (positive)
• Show the electrolyte and name the ions
• Draw arrows showing ion movement (cations to cathode, anions to anode)
• Write the half-equations at the correct electrode
• Note the observation at each electrode (gas bubbles, metal deposited, colour change)

Drawing from memory is a powerful test. If you can reproduce the full labelled diagram for the electrolysis of concentrated aqueous sodium chloride without notes, you know the topic. See our diagram guide for examiner-standard labelling.

Use a comparison table

Feature	Molten PbBr2	Aqueous CuSO4	Aqueous NaCl (conc.)
Cathode product	Lead	Copper	Hydrogen
Anode product	Bromine	Oxygen	Chlorine
Why cathode product?	Only cation present	Cu below H in reactivity series	Na above H in reactivity series
Why anode product?	Only anion present	SO4 2- is not a halide	Cl- is a halide

Building this table yourself, then recreating it from memory, tests both the facts and the reasoning.

Link to the reactivity series

The reactivity series is the key that unlocks electrode product predictions. If you know the series well, you can predict the cathode product for any aqueous electrolysis without memorising individual examples. This is a case where understanding beats memorising: learn the principle, and it works for every question.

Simple cells and electrochemical cells

A simple cell produces electricity from two different metals in an electrolyte. The more reactive metal is the negative terminal (it loses electrons more readily), and the less reactive metal is the positive terminal.

Key points for 0620:

• The voltage depends on how far apart the metals are in the reactivity series (Zn/Cu gives a higher voltage than Fe/Cu)
• The more reactive metal corrodes (is oxidised) — it gets smaller
• The less reactive metal is the cathode (site of reduction)
• The electrolyte must contain ions to complete the circuit

Practise predicting which metal is the negative terminal and which will corrode for different metal pairs. The reactivity series gives you the answer every time.

Hydrogen fuel cells

Extended students need to know how hydrogen fuel cells work:

• Hydrogen is oxidised at one electrode: 2H2 → 4H+ + 4e-
• Oxygen is reduced at the other: O2 + 4H+ + 4e- → 2H2O
• The overall reaction is 2H2 + O2 → 2H2O
• The only product is water, making it a clean energy source
• Advantages over fossil fuels: no CO2 emissions, no pollutant gases
• Disadvantages: hydrogen is difficult to store (it is a gas, needs high pressure), production often uses fossil fuels, infrastructure is limited

Common exam errors

• Confusing the polarity of electrodes between electrolysis and simple cells
• Forgetting to balance half-equations (electrons must balance with ion charges)
• Writing “copper dissolves at the cathode” when copper is deposited at the cathode (copper dissolves at a copper anode)
• Stating that “electricity flows through the solution” — ions carry the current in the solution; electrons flow in the external circuit only
• Forgetting observations: bubbles of gas, colour changes in solution, metal deposited

For exam-specific strategies, see the electrochemistry exam guide and the electrolysis exam technique article.