State that metal objects are electroplated to improve their appearance and resistance to corrosion

Electrochemistry – Electrolysis (IGCSE 0620)

1. What is electrolysis?

Electrolysis is the use of electrical energy to drive a non‑spontaneous redox reaction. It occurs in an electrolytic cell where an external power source forces ions to move toward the electrodes.

2. Main components of an electrolytic cell

3. Half‑reactions required by the Cambridge syllabus (Core 4.1)

General rule (Core 4.1)

• Metals (or H⁺ ions) are reduced at the cathode – a metal is deposited or hydrogen gas is formed.
• Non‑metals (other than hydrogen) are oxidised at the anode – the elemental non‑metal is liberated.

Predicting products for any binary molten compound

1. If the compound contains a metal + non‑metal, the metal cation is reduced at the cathode (metal deposited) and the non‑metal anion is oxidised at the anode (gives the elemental non‑metal).
2. If both ions are non‑metals, compare standard reduction potentials; the ion with the higher (more positive) potential is reduced at the cathode, the other is oxidised at the anode.

Worked example – molten sodium sulphide (Na₂S)

• Ionisation: Na₂S(l) → 2 Na⁺ + S²⁻
• At the cathode: Na⁺ + e⁻ → Na (s) (reduction of the metal cation)
• At the anode: S²⁻ → S (s) + 2 e⁻ (oxidation of the non‑metal anion)
• Overall: Na₂S(l) → 2 Na (s) + S (s)

4. Safety considerations (Core 4.1)

• Wear safety goggles and gloves; some electrolytes (e.g. H₂SO₄) are corrosive.
• Carry out the experiment in a well‑ventilated area – gases such as H₂, Cl₂ or O₂ may be produced.
• Never touch the electrodes while the cell is operating; they can become hot.
• Use a suitable power supply with a current‑limiting feature to avoid overheating.

5. Hydrogen–oxygen fuel cell (Supplement 4)

A fuel cell converts the chemical energy of hydrogen and oxygen directly into electrical energy.

• Overall reaction: 2 H₂ + O₂ → 2 H₂O ΔE° ≈ 1.23 V
• At the anode: H₂ → 2 H⁺ + 2 e⁻ (oxidation)
• At the cathode: ½ O₂ + 2 H⁺ + 2 e⁻ → H₂O (reduction)
• The flow of electrons through an external circuit supplies usable electricity – a clean, renewable‑energy technology.

6. Electroplating – definition

Electroplating is a specialised type of electrolysis in which a thin, uniform layer of metal is deposited onto the surface of a solid object (the substrate) to improve its appearance, corrosion resistance, wear resistance or to restore dimensions.

7. How electroplating works (Core 4.2)

1. Set‑up – The object to be plated is connected as the cathode. The metal that will form the coating is placed in the electrolyte as a soluble salt (e.g. CuSO₄) and usually also serves as the anode.
2. Ion migration – When the power supply is switched on, metal ions move toward the cathode.
3. Reduction at the cathode – Metal ions gain electrons and deposit as a solid layer:
   Mⁿ⁺ + n e⁻ → M (s)
4. Oxidation at the anode – The anode metal dissolves, replenishing the ion concentration:
   M (s) → Mⁿ⁺ + n e⁻
5. Result – A continuous, adherent metallic coating builds up on the substrate.

8. Reasons for electroplating (Core 4.2)

• Improved appearance – shiny, decorative surfaces (e.g. silver‑plated jewellery).
• Increased corrosion resistance – a protective barrier prevents the underlying metal from reacting with the environment (e.g. zinc‑plated nails).
• Enhanced wear resistance – hard metals such as chromium reduce abrasion on tools and fittings.
• Restoration of dimensions – worn parts can be rebuilt by adding material.

9. Common plating metals and their typical uses

10. Example: Copper plating of an iron nail (practical activity)

Set‑up

• Electrolyte: 0.5 M CuSO₄ solution (aqueous).
• Anode: clean copper sheet.
• Cathode: iron nail (the object to be plated), attached with a crocodile clip.
• Power supply: DC source, 2 V, current ≈ 0.5 A.

Observations

• Cu²⁺ ions migrate to the nail and are reduced: Cu²⁺ + 2 e⁻ → Cu(s). A thin reddish‑gold coating appears on the nail.
• At the copper anode: Cu(s) → Cu²⁺ + 2 e⁻, maintaining the ion concentration.
• The plated nail shows a brighter surface and is less prone to rusting when exposed to moist air.

11. Summary

• Electrolysis uses electricity to force a non‑spontaneous redox reaction in an electrolytic cell.
• Key components: power supply, electrolyte, anode (positive, oxidation), cathode (negative, reduction), and, when required, inert electrodes.
• Core 4.1 requires knowledge of the half‑reactions for molten PbBr₂, concentrated NaCl, and dilute H₂SO₄, plus the rule that metals (or H⁺) are reduced at the cathode and non‑metals are oxidised at the anode.
• Predicting products for any binary molten compound follows a simple decision‑tree; a worked example (Na₂S) demonstrates the method.
• Safety: wear goggles, work in a ventilated area, and be aware of hazardous gases.
• Supplement 4 introduces the hydrogen‑oxygen fuel cell, showing how redox reactions can generate electricity directly.
• Electroplating applies the same principles to deposit a thin metal layer, providing aesthetic, protective, and mechanical benefits. Common plating metals (Cu, Ni, Cr, Ag, Zn) each give specific advantages.