Electrolysis, Electroplating and Fuel Cells

CIE IGCSE Chemistry· 4.1, 4.2· 45 min read

1. Core: Electrolysis Fundamentals★★☆☆☆⏱ 10 min

Electrolysis is the process of breaking down (decomposing) ionic compounds using an electric current. It requires three core components: a power source, two electrodes (usually inert graphite or platinum for Core syllabi), and an electrolyte (molten or aqueous ionic compound with mobile ions to carry charge).

📘 Definition

Electrolysis

The decomposition of an ionic compound, when molten or dissolved in water, by passing an electric current through it.

Example:

Electrolysis of molten sodium chloride breaks it into sodium metal and chlorine gas.

Ions move to oppositely charged electrodes during electrolysis: negative anions move to the positive anode (where oxidation, loss of electrons, occurs) and positive cations move to the negative cathode (where reduction, gain of electrons, occurs).

📐 Worked Example

Name the products formed at the anode and cathode when molten lead(II) bromide is electrolysed using inert graphite electrodes.

1
1. Identify ions present in the molten electrolyte: Pb²+ (cation) and Br⁻ (anion, no water ions present as it is molten)
2
1. Cations are attracted to the negative cathode: Pb²+ gains electrons to form solid lead metal
3
1. Anions are attracted to the positive anode: Br⁻ loses electrons to form bromine gas
4
1. Final products: Cathode = lead metal, Anode = bromine gas

Exam tip:

Always confirm if the electrolyte is molten or aqueous first, as product prediction rules change completely for aqueous solutions.

2. Core: Electrolysis of Aqueous Solutions★★★☆☆⏱ 10 min

Aqueous electrolytes contain H+ and OH⁻ ions from water, alongside the ions from the dissolved ionic compound. The discharge series ranks ions by ease of release at each electrode to predict products.

📘 Definition

Discharge Series (Aqueous Electrolysis)

Ranking of ions by ease of discharge: At the cathode, less reactive metals (below hydrogen in the reactivity series) discharge before H+ ions, which form hydrogen gas if the metal is more reactive. At the anode, concentrated halide ions (Cl⁻, Br⁻, I⁻) discharge before OH⁻ ions, which form oxygen gas if halides are dilute or absent.

📐 Worked Example

Predict the products at each electrode when dilute copper(II) sulfate solution is electrolysed using inert graphite electrodes.

1
1. List all ions present: Cu²+, SO₄²⁻, H+, OH⁻
2
1. Cathode rule: Copper is below hydrogen in the reactivity series, so Cu²+ discharges to form solid copper metal
3
1. Anode rule: No concentrated halide ions are present, so OH⁻ discharges to form oxygen gas and water
4
1. Final products: Cathode = copper metal, Anode = oxygen gas

Exam tip:

If the anode is made of a reactive metal (e.g. copper instead of graphite), it will dissolve during electrolysis to replace metal ions in the solution, instead of anions discharging.

3. Core: Electroplating★★☆☆☆⏱ 8 min

Electroplating is an industrial electrolytic process used to coat cheaper, more reactive metals (e.g. iron, steel) with a thin layer of a more expensive, corrosion-resistant metal (e.g. silver, chromium, nickel) to improve appearance or prevent rusting.

📘 Definition

Electroplating Setup

Three required components: 1. Cathode = object to be plated, 2. Anode = pure bar of the coating metal, 3. Electrolyte = aqueous salt solution of the coating metal.

Example:

Plating a steel spoon with silver: Spoon = cathode, pure silver bar = anode, silver nitrate solution = electrolyte.

📐 Worked Example

A student wants to electroplate an iron key with nickel. List the required components for the electrolytic setup.

1
1. Cathode: Iron key (negative terminal, attracts Ni²+ ions to coat its surface)
2
1. Anode: Pure nickel bar (positive terminal, dissolves during electrolysis to replace Ni²+ ions used up in solution)
3
1. Electrolyte: Aqueous nickel(II) sulfate solution (provides mobile Ni²+ ions for plating)

Exam tip:

Never mix up anode and cathode for electroplating: the object you want to plate always connects to the negative terminal (cathode).

4. Core: Hydrogen-Oxygen Fuel Cells★★★☆☆⏱ 7 min

Fuel cells are electrochemical cells that convert the chemical energy of a fuel (usually hydrogen) and oxygen directly into electrical energy, with no combustion step. They have much higher energy efficiency than fossil fuel power stations or internal combustion engines.

📘 Definition

Hydrogen-Oxygen Fuel Cell

A cell that uses the redox reaction between hydrogen and oxygen to produce electricity, with water as the only waste product. Key advantages: No greenhouse gas or toxic emissions, quiet operation, high energy density for transport. Key disadvantages: Hydrogen is highly flammable, expensive to produce and store, limited refuelling infrastructure.

📐 Worked Example

State two advantages of using hydrogen fuel cells instead of petrol internal combustion engines in passenger cars.

1
1. First advantage: The only waste product is water, so no carbon dioxide, nitrogen oxides or carbon monoxide emissions that contribute to climate change or air pollution.
2
1. Second advantage: Fuel cells have ~60% energy conversion efficiency, compared to ~20% efficiency for internal combustion engines, so less energy is wasted as heat.

5. Extended Only: Half Equations for Electrode Reactions★★★★☆Extended only⏱ 10 min

Extended candidates are required to write balanced half equations for reactions at each electrode, showing electron transfer. Oxidation reactions (at the anode) have electrons on the product side (lost), while reduction reactions (at the cathode) have electrons on the reactant side (gained).

📘 Definition

Half Equation (Extended)

An ionic equation that shows either the oxidation or reduction reaction taking place at a single electrode, including electron transfer.

Example:

Reduction of copper ions at the cathode: $C u^{2 +} + 2 e^{-} \to C u$

📐 Worked Example

Write the balanced half equation for the reaction at the anode when concentrated sodium chloride solution (brine) is electrolysed using inert electrodes.

1
1. Identify the discharged ion: Concentrated Cl⁻ ions discharge before OH⁻ ions at the anode
2
1. Oxidation occurs at the anode, so each Cl⁻ loses 1 electron to form a Cl atom
3
1. Chlorine is diatomic, so two Cl atoms combine to form Cl₂ gas
4
$2Cl^- \rightarrow Cl_2 + 2e^-$
5
1. Check balance: 2 Cl atoms on each side, total charge of -2 on both sides of the equation

Exam tip:

Always check that the total charge on both sides of a half equation is equal, and that the number of atoms of each element is balanced.

6. Common Pitfalls

Wrong move:

Assuming solid ionic compounds can be electrolysed

Why:

Ions are fixed in a solid ionic lattice and cannot move to carry charge

Correct move:

Only electrolyse ionic compounds when molten or dissolved in water, so ions are mobile

Wrong move:

Applying the same product rules to molten and aqueous electrolytes

Why:

Aqueous electrolytes contain H+ and OH⁻ ions from water that are not present in molten electrolytes

Correct move:

First confirm if the electrolyte is molten or aqueous before applying discharge rules

Wrong move:

Setting the object to be plated as the anode in electroplating setups

Why:

Positive metal ions are attracted to the negative cathode, so only the cathode gets coated

Correct move:

Always set the object to be plated as the negative cathode, and the coating metal as the positive anode

Wrong move:

Forgetting reactive anodes dissolve during electrolysis

Why:

Reactive metals at the anode lose electrons more easily than anions in solution, so they oxidise and dissolve

Correct move:

If the anode is not inert (e.g. copper instead of graphite), it will dissolve to release metal ions into the electrolyte

Wrong move:

Writing half equations with electrons on the wrong side

Why:

Oxidation (anode) loses electrons so electrons are on the product side; reduction (cathode) gains electrons so electrons are on the reactant side

Correct move:

Label the electrode first to confirm if oxidation or reduction is taking place, then place electrons accordingly

Wrong move:

Stating fuel cells burn hydrogen to produce electricity

Why:

Fuel cells do not involve combustion; they directly convert chemical energy to electrical energy via redox reactions

Correct move:

Describe fuel cell operation as a direct electrochemical reaction between hydrogen and oxygen, with no combustion step

7. Quick Reference Cheatsheet

Concept	Core Rule	Extended Rule
Electrolyte requirement	Must be molten/aqueous ionic compound with mobile ions	Same as Core + ionic charge determines number of electrons transferred per ion
Cathode product (molten)	Metal from the ionic compound	Same as Core + write reduction half-equation
Anode product (molten)	Non-metal from the ionic compound	Same as Core + write oxidation half-equation
Cathode product (aqueous)	Hydrogen gas if metal > H in reactivity series, else metal	Same as Core + balance half-equation for electron count
Anode product (aqueous)	Halogen if concentrated halide present, else oxygen gas	Same as Core + account for diatomic non-metals in half-equations
Electroplating setup	Object = cathode, coating metal = anode, electrolyte = coating metal salt solution	Same as Core + construct ionic half-equations for the electrode reactions (metal dissolves at the anode, metal deposits at the cathode)
Hydrogen fuel cell	Waste = water, advantage = no CO₂ emissions, disadvantage = H₂ storage issues	Same as Core + write overall and half-equations for cell reaction

8. Frequently Asked

Can solid ionic compounds be electrolysed?

No. Solid ionic compounds have ions fixed in a lattice, so no mobile charge carriers to conduct electricity. Only molten or aqueous ionic compounds can be electrolysed.

Why is hydrogen gas produced at the cathode in some aqueous electrolysis setups?

If the metal cation is more reactive than hydrogen (e.g. sodium, magnesium) in the reactivity series, H+ ions from water are discharged in place of the metal cation to form hydrogen gas.

Going deeper

syllabusCIE IGCSE Chemistry 0620 2026-2028 SyllabusOfficial syllabus for reference
practiceElectrochemistry Structured Practice QuestionsNo MCQ, aligned to Core/Extended tiers

What's Next

Now that you have mastered core and extended content for electrolysis, electroplating and fuel cells, you are ready to apply this knowledge to other electrochemistry topics in the CIE IGCSE Chemistry 0620 syllabus. Next, you should practice predicting electrolysis products for a range of molten and aqueous setups, and work through extended response questions on electroplating applications and fuel cell comparisons with fossil fuel power sources. If you are sitting the Extended paper, make sure you practice writing balanced half-equations for all electrode reactions, as this is a common high-mark question in Paper 4. You should also link this content to redox reaction definitions you learned earlier, as oxidation and reduction are core to all electrolytic processes.