Electrolysis

Definition of Electrolysis

• Electrolysis is the decomposition of an ionic compound, when molten or in aqueous solution, by the passage of an electric current.
• The process requires:
  • An electrolyte: molten ionic compound or aqueous ionic solution containing free-moving ions.
  • Electrodes: conductors where oxidation and reduction occur.
  • A power supply: usually a direct current (DC) source.

Basic Components of an Electrolytic Cell

• Anode:
  • Positive electrode connected to the positive terminal of the power supply.
  • Site of oxidation (loss of electrons).
  • Attracts anions (negative ions).
• Cathode:
  • Negative electrode connected to the negative terminal of the power supply.
  • Site of reduction (gain of electrons).
  • Attracts cations (positive ions).
• Electrolyte:
  • Molten or aqueous ionic substance that conducts electricity by movement of ions.
  • In molten state: contains only the compound’s ions.
  • In aqueous state: contains compound’s ions and ions from water.

Movement of Charge

• Electrons flow through the external circuit from the anode to the cathode.
• In the electrolyte:
  • Cations move toward the cathode to gain electrons (reduction).
  • Anions move toward the anode to lose electrons (oxidation).
• The overall process involves:
  • Electron transfer at the electrodes.
  • Ion migration in the electrolyte.

Electrolysis Examples and Observations

1. Molten Lead(II) Bromide (PbBr₂)

• Electrolyte: molten PbBr₂.
• Cathode (–): Pb²⁺ ions gain electrons → lead metal deposited.
  • Pb²⁺ + 2e⁻ → Pb (reduction).
  • Observation: grey molten lead forms at the cathode.
• Anode (+): Br⁻ ions lose electrons → bromine gas.
  • 2Br⁻ → Br₂ + 2e⁻ (oxidation).
  • Observation: red-brown bromine vapour released.

2. Concentrated Aqueous Sodium Chloride (brine)

• Electrolyte: concentrated NaCl(aq).
• Cathode (–): H⁺ ions from water reduced to hydrogen gas.
  • 2H⁺ + 2e⁻ → H₂.
  • Observation: colourless gas, pops with lit splint.
• Anode (+): Cl⁻ ions oxidised to chlorine gas.
  • 2Cl⁻ → Cl₂ + 2e⁻.
  • Observation: pale green gas with choking smell.
• Sodium hydroxide remains in solution.

3. Dilute Sulfuric Acid (H₂SO₄)

• Electrolyte: dilute H₂SO₄.
• Cathode (–): H⁺ ions reduced to hydrogen gas.
  • 2H⁺ + 2e⁻ → H₂.
• Anode (+): OH⁻ ions from water oxidised to oxygen gas.
  • 4OH⁻ → O₂ + 2H₂O + 4e⁻.
• Observation: twice as much gas at cathode as anode (2:1 ratio H₂:O₂).

4. Aqueous Copper(II) Sulfate (CuSO₄) with Inert Electrodes

• Electrolyte: CuSO₄(aq).
• Cathode (–): Cu²⁺ ions reduced to copper metal.
  • Cu²⁺ + 2e⁻ → Cu.
  • Observation: reddish-brown copper coating.
• Anode (+): OH⁻ ions oxidised to oxygen gas.
  • 4OH⁻ → O₂ + 2H₂O + 4e⁻.

5. Aqueous Copper(II) Sulfate with Copper Electrodes

• Cathode (–): Cu²⁺ ions gain electrons → copper deposited.
• Anode (+): Copper atoms lose electrons → Cu²⁺ ions enter solution.
• Reaction:
  • Anode: Cu → Cu²⁺ + 2e⁻.
  • Cathode: Cu²⁺ + 2e⁻ → Cu.
• Mass of cathode increases; mass of anode decreases.
• Used in purification of copper.

General Rules for Electrolysis Products

• Molten Binary Compounds:
  • Cathode: metal formed.
  • Anode: non-metal formed.
• Aqueous Solutions:
  • Cathode: H₂ forms if metal is more reactive than hydrogen; otherwise metal forms.
  • Anode: halogen gas forms if halide ions present; otherwise oxygen forms from OH⁻.

Half-equations

• Oxidation at anode: X⁻ → X + e⁻ (loss of electrons).
• Reduction at cathode: Y⁺ + e⁻ → Y (gain of electrons).

Electroplating

• Purpose: improve appearance and corrosion resistance.
• Method:
  • Cathode: object to be plated.
  • Anode: metal to be plated onto object.
  • Electrolyte: aqueous solution of metal salt.
• Example: silver plating using silver anode and AgNO₃(aq).