OxfordAQA International GCSE Chemistry: Electrolysis (9202) – Revision Notes

These revision notes condense section 3.3.2 Electrolysis of the OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards. They cover statements a to j and the required practical on copper sulfate solution. The qualification is untiered, so all of it is for every candidate. The electrolysis study guide has the full teaching and worked examples.

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Key definitions (3.3.2 a to e)

Term	Meaning
Electrolysis	Breaking down an ionic substance, molten or in solution, by passing an electric current through it
Electrolyte	The molten or dissolved ionic substance that is broken down
Cathode	Negative electrode; positive ions move to it
Anode	Positive electrode; negative ions move to it
Oxidation	Loss of electrons (happens at the anode)
Reduction	Gain of electrons (happens at the cathode)

Why melt or dissolve? In a solid ionic lattice the ions are fixed in place. Melting or dissolving frees them to move, so they can carry charge through the liquid.

Molten compound with two ions: the metal forms at the cathode, the non-metal at the anode. Both products are elements.

Electron flow: electrons move through the external wires. Inside the electrolyte the current is carried by moving ions.

Half equations (3.3.2 f)

Method in steps:

1. Ion on the left, element on the right.
2. Balance atoms. Halogens and oxygen are diatomic: Cl₂, Br₂, I₂, O₂.
3. Add electrons to balance charge. Cathode: electrons on the left. Anode: electrons on the right (or “– e⁻” on the left).

Electrode	Ion	Half equation
Cathode	Na⁺	Na⁺ + e⁻ → Na
Cathode	Cu²⁺	Cu²⁺ + 2e⁻ → Cu
Cathode	Ag⁺	Ag⁺ + e⁻ → Ag
Cathode	Al³⁺	Al³⁺ + 3e⁻ → Al
Cathode	H⁺	2H⁺ + 2e⁻ → H₂
Anode	Cl⁻	2Cl⁻ → Cl₂ + 2e⁻ (or 2Cl⁻ – 2e⁻ → Cl₂)
Anode	Br⁻	2Br⁻ → Br₂ + 2e⁻
Anode	O²⁻ (molten oxide)	2O²⁻ → O₂ + 4e⁻
Anode	OH⁻ (from water)	4OH⁻ → O₂ + 2H₂O + 4e⁻

Quick check: total charge on the left must equal total charge on the right. For 2O²⁻ → O₂ + 4e⁻: left is 4–, right is 0 + 4– = 4–.

Worked reminders

Molten lithium oxide. Ions: Li⁺ and O²⁻. Cathode: Li⁺ + e⁻ → Li. Anode: 2O²⁻ → O₂ + 4e⁻. Every O₂ molecule releases four electrons, and each Li⁺ takes one, so four lithium atoms form for every oxygen molecule.

Spot the process from the equation. Electrons on the left of the arrow are gained, so the change is reduction. Electrons on the right (or subtracted on the left) are lost, so the change is oxidation.

Predicting aqueous products in three questions:

1. Which ions are present? Add H⁺ and OH⁻ from water to the ions of the salt.
2. Cathode: is the metal more or less reactive than hydrogen?
3. Anode: is there a halide, and is it concentrated?

Mixtures of ions in solution (3.3.2 g)

Water adds H⁺ and OH⁻ ions to every aqueous electrolyte.

Cathode rule (reactivity):

• metal more reactive than hydrogen → hydrogen
• metal less reactive than hydrogen (copper, silver) → the metal

Anode rule (concentration matters):

• concentrated chloride, bromide or iodide → the halogen
• dilute halide, or sulfate or nitrate → oxygen

Solution (carbon electrodes)	Cathode	Anode
Copper sulfate	Copper	Oxygen
Sodium chloride (concentrated)	Hydrogen	Chlorine
Potassium sulfate	Hydrogen	Oxygen
Silver nitrate	Silver	Oxygen
Concentrated sodium bromide	Hydrogen	Bromine

Required practical: copper sulfate solution

Set-up: two clean carbon rods in copper sulfate solution, connected to a low-voltage direct-current supply. Weigh the cathode before and after.

What you see and why:

• Cathode: pink-brown copper coating; cathode mass increases. Cu²⁺ + 2e⁻ → Cu.
• Anode: colourless gas bubbles that relight a glowing splint, so oxygen. 4OH⁻ → O₂ + 2H₂O + 4e⁻.
• Solution: blue colour fades as Cu²⁺ ions are used up.

Practical points:

• Rinse the cathode gently (rough handling knocks off loose copper) and dry it before reweighing (water left on it adds to the mass).
• Keep the rods apart so they do not touch.
• Eye protection: copper sulfate solution is an irritant.
• To spot an anomaly, compare each reading with the steady trend in mass gain.

Improving and evaluating the practical:

• Repeat the experiment and take a mean of the mass gains.
• Keep the current steady, using an ammeter and a variable resistor if available.
• Clean the electrodes with emery paper first so the copper sticks evenly.
• A conclusion should link each observation to a product: coating means copper, relit splint means oxygen.

Electroplating (3.3.2 h)

Part	What it is
Cathode	The object to be plated
Anode	The plating metal (copper or silver)
Electrolyte	A solution containing ions of the plating metal

• Anode: metal atoms lose electrons and dissolve (Cu → Cu²⁺ + 2e⁻; Ag → Ag⁺ + e⁻).
• Cathode: metal ions gain electrons and coat the object.
• The concentration of the electrolyte stays roughly constant.
• Reasons: appearance, durability, prevention of corrosion.

Aluminium (3.3.2 i)

• Electrolyte: aluminium oxide dissolved in molten cryolite.
• Why cryolite: aluminium oxide alone melts at over 2000 °C. The mixture is molten at a much lower temperature, so less energy is needed and costs fall.
• Negative electrode: Al³⁺ + 3e⁻ → Al (aluminium forms).
• Positive electrode (carbon): 2O²⁻ → O₂ + 4e⁻ (oxygen forms).
• The carbon anodes react with oxygen, C + O₂ → CO₂, so they wear away and are replaced.
• Overall: 2Al₂O₃ → 4Al + 3O₂.
• Large amounts of energy (electricity for melting and for decomposing the oxide) make aluminium expensive.

Reacting-mass shortcut. From 2Al₂O₃ → 4Al + 3O₂, one mole of aluminium oxide (102 g) gives two moles of aluminium (54 g). So the maximum mass of aluminium is always 54/102 of the mass of oxide used. Show the mole steps in an answer anyway; the quantitative chemistry revision notes recap them.

Why the cell is expensive to run: electricity is needed both to keep the mixture molten and to decompose the oxide, and the carbon anodes are used up and must be replaced.

Sodium chloride solution (3.3.2 j)

Where	Product	Reason
Cathode	Hydrogen	Sodium is more reactive than hydrogen, so H⁺ is discharged
Anode	Chlorine	Chloride ions are at high concentration
Solution	Sodium hydroxide	Na⁺ and OH⁻ ions remain

Uses: sodium hydroxide for soap; chlorine for bleach and plastics.

Must-know distinctions

• Cathode vs anode: cathode is negative and attracts cations; anode is positive and attracts anions.
• Oxidation vs reduction: loss vs gain of electrons. OIL RIG.
• Molten vs aqueous: molten gives the metal and non-metal of the compound; aqueous may give hydrogen and/or oxygen from water instead.
• Ions vs electrons: charge travels through the electrolyte as moving ions but through the wires as electrons.
• Reactivity vs concentration: reactivity decides the cathode product; concentration also matters at the anode.

Quick self-test

1. Why does solid potassium chloride not conduct electricity?
2. Name the products when molten potassium chloride is electrolysed.
3. Write the cathode half equation for molten potassium chloride.
4. Complete: …Br⁻ → Br₂ + …
5. Is the reaction Fe³⁺ + 3e⁻ → Fe oxidation or reduction? Explain.
6. Predict the products at each electrode for concentrated copper(II) bromide solution.
7. Predict the products at each electrode for potassium nitrate solution.
8. In the copper sulfate practical, how would you show the anode gas is oxygen?
9. A brass key is to be copper plated. Name the anode and the electrolyte.
10. Give one reason, other than appearance, for electroplating.
11. Why are the positive electrodes in aluminium manufacture replaced regularly?
12. Calculate the maximum mass of aluminium from 3.40 kg of aluminium oxide. (Aᵣ: Al = 27, O = 16)

Answers

1. In the solid lattice each ion stays locked in place, so no charged particles can travel.
2. Potassium at the cathode; chlorine at the anode.
3. K⁺ + e⁻ → K
4. 2Br⁻ → Br₂ + 2e⁻
5. Reduction, because the iron ion gains electrons.
6. Cathode: copper (less reactive than hydrogen). Anode: bromine (bromide at high concentration).
7. Cathode: hydrogen (potassium more reactive than hydrogen). Anode: oxygen (nitrate is not a halide).
8. Collect it in an inverted test tube and insert a glowing splint; it relights.
9. Anode: copper. Electrolyte: copper sulfate solution (a solution of copper ions).
10. Durability, or prevention of corrosion.
11. The carbon reacts with the oxygen formed to make carbon dioxide, so the electrodes wear away.
12. Mᵣ Al₂O₃ = 102; 3.40 ÷ 102 = 0.0333 kmol; Al = 2 × 0.0333 = 0.0667 kmol; mass = 0.0667 × 27 = 1.80 kg.

Where marks are usually lost

• Writing single atoms (Cl, O, Br) where the molecule (Cl₂, O₂, Br₂) is needed.
• Unbalanced charge in a half equation, for example 2O²⁻ → O₂ + 2e⁻.
• Saying the ions “are free to move” in the solid, or that electrons carry the current through the electrolyte.
• Naming sodium (or potassium) as a product at the cathode in an aqueous solution.
• Forgetting that a dilute halide solution gives oxygen at the anode.
• In electroplating, making the object the anode, or using an electrolyte that does not contain ions of the plating metal.
• Describing cryolite as a catalyst instead of saying it lowers the melting temperature and so saves energy.
• Saying the carbon anode “dissolves” instead of reacting with oxygen to form carbon dioxide.
• In the brine answer, explaining the cathode product without referring to reactivity.

Official syllabus

OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards, published by OxfordAQA. These notes cover section 3.3.2 Electrolysis and its required practical.