Definition: Electrolysis is the decomposition of an ionic compound when molten or in aqueous solution by the passage of an electric current.

Why Only Ionic Compounds?

• Solid state: Ions are fixed in place in the lattice, cannot move → no conductivity
• Molten or dissolved: Ions are free to move and carry charge → conducts electricity

Key Terms:

表 TermDefinitionWhat it attracts 表 ElectrodeRod (metal or graphite) that conducts electricity into/out of the electrolyte- AnodePositive electrodeAnions (negative ions) CathodeNegative electrodeCations (positive ions) ElectrolyteIonic compound in molten or dissolved solution that conducts electricity- AnionNegatively charged ionAttracted to anode CationPositively charged ionAttracted to cathode 表

(a) Molten Lead(II) Bromide (PbBr₂)

表 ElectrodeIonProductObservation 表 英语Cathode (−)Pb²⁺Grey lead metalLead deposits on electrode Anode (+)Br⁻Brown bromine gasBubbling at anode, pungent smell 表

(b) Concentrated Aqueous Sodium Chloride (Brine)

表 ElectrodeIon DischargedProductObservation 表 英语Cathode (−)H⁺ (from water)Hydrogen gasBubbling Anode (+)Cl⁻Chlorine gasBubbling, pungent smell, bleaches damp litmus 表

(c) Dilute Sulfuric Acid (H₂SO₄)

表 ElectrodeIon DischargedProductObservation 表 英语Cathode (−)H⁺Hydrogen gasBubbling, squeaky pop test Anode (+)OH⁻ (from water)Oxygen gasBubbling, relights glowing splint 表

At the Cathode (Negative Electrode):

• Metals are produced (if metal ions are discharged)
• Hydrogen is produced (if H⁺ ions are discharged)

At the Anode (Positive Electrode):

• Non-metals (other than hydrogen) are produced
• Examples: Chlorine (Cl₂), Bromine (Br₂), Iodine (I₂), Oxygen (O₂)

Rule: For molten binary compounds (two elements), the products are always the elements themselves.

Worked Example 1: Molten Potassium Chloride (KCl)

• Cathode: K⁺ + e⁻ → K (potassium metal)
• Anode: 2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas)

Worked Example 2: Molten Magnesium Oxide (MgO)

• Cathode: Mg²⁺ + 2e⁻ → Mg (magnesium metal)
• Anode: 2O²⁻ → O₂ + 4e⁻ (oxygen gas)

What is Electroplating? Coating one metal with a layer of another metal using electrolysis.

Reasons for Electroplating:

• Corrosion protection: Prevents rusting/oxidation
• Improved appearance: Makes objects look more attractive
• Increased durability: Adds a hard, wear-resistant layer

Examples: Galvanising (zinc on steel), chromium plating, silver plating, gold plating

Setup for Electroplating:

• Cathode (−): Object to be plated
• Anode (+): Pure metal to be plated
• Electrolyte: Aqueous salt of the plating metal

Example: Plating Iron with Tin

• Cathode: Iron strip (object to be plated)
• Anode: Tin metal
• Electrolyte: Tin(II) chloride solution (SnCl₂)

Half-Equations:

• Cathode: Sn²⁺(aq) + 2e⁻ → Sn(s) (tin deposited on iron)
• Anode: Sn(s) → Sn²⁺(aq) + 2e⁻ (tin dissolves)

How Charge Flows in Electrolysis:

(a) External Circuit: Electrons flow from the negative terminal of the power supply to the cathode, then from the anode back to the positive terminal.

(b) At the Electrodes:

• Cathode (−): Cations gain electrons → reduction
• Anode (+): Anions lose electrons → oxidation

(c) In the Electrolyte: Ions move and carry charge

• Cations (+) move towards the cathode
• Anions (−) move towards the anode

Using Inert Electrodes (Graphite/Platinum):

表 ElectrodeIon DischargedProductObservation 表 英语Cathode (−)Cu²⁺Copper metalPink/brown deposit on electrode Anode (+)OH⁻ (from water)Oxygen gasBubbling, relights glowing splint 表

Half-Equations: Cathode: Cu²⁺(aq) + 2e⁻ → Cu(s) | Anode: 4OH⁻(aq) → O₂(g) + 2H₂O(l) + 4e⁻

Using Copper Electrodes (Active Electrodes):

表 ElectrodeProcessObservation 表 英语Cathode (−)Cu²⁺ reduced to CuGains mass (copper deposited) Anode (+)Cu oxidised to Cu²⁺Loses mass (copper dissolves) 表

Rule for Anode Products: The concentration of the solution affects which ion is discharged at the anode.

表 SolutionIons PresentProduct at AnodeReason 表 英语Concentrated halideCl⁻, Br⁻, I⁻Halogen gas (Cl₂, Br₂, I₂)Halide ions are discharged Dilute halideCl⁻, Br⁻, I⁻ + OH⁻Oxygen gasOH⁻ is discharged instead 表

Examples:

• Concentrated NaCl: Cl⁻ discharged → chlorine gas at anode
• Dilute NaCl: OH⁻ discharged → oxygen gas at anode

Common Half-Equations:

表 ProductHalf-EquationType 表 英语Hydrogen2H⁺ + 2e⁻ → H₂Reduction (cathode) Metal (e.g., Cu)Cu²⁺ + 2e⁻ → CuReduction (cathode) Chlorine2Cl⁻ → Cl₂ + 2e⁻Oxidation (anode) Bromine2Br⁻ → Br₂ + 2e⁻Oxidation (anode) Iodine2I⁻ → I₂ + 2e⁻Oxidation (anode) Oxygen (from OH⁻)4OH⁻ → O₂ + 2H₂O + 4e⁻Oxidation (anode) 表

Worked Example: Molten Lead(II) Bromide

• Cathode: Pb²⁺ + 2e⁻ → Pb(l)
• Anode: 2Br⁻ → Br₂(g) + 2e⁻

Worked Example: Concentrated Sodium Chloride

• Cathode: 2H⁺ + 2e⁻ → H₂(g)
• Anode: 2Cl⁻ → Cl₂(g) + 2e⁻

What is a Fuel Cell? An electrochemical cell that converts chemical energy from a fuel into electrical energy.

Hydrogen-Oxygen Fuel Cell:

• Fuel: Hydrogen (H₂)
• Oxidant: Oxygen (O₂)
• Only product: Water (H₂O)

Overall Reaction: 2H₂(g) + O₂(g) → 2H₂O(l)

Half-Equations:

• Anode (−): H₂(g) → 2H⁺(aq) + 2e⁻
• Cathode (+): O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l)

Advantages of Hydrogen Fuel Cells:

• ✅ Only product is water (no pollution, no CO₂)
• ✅ Renewable (hydrogen can be produced from water)
• ✅ Higher energy per kilogram than petrol/diesel
• ✅ No moving parts → no power loss in transmission
• ✅ Quieter than petrol engines (less noise pollution)

Disadvantages of Hydrogen Fuel Cells:

• ❌ Hydrogen production may use fossil fuels (releases CO₂)
• ❌ Electrolysis of water requires large amounts of electricity
• ❌ Expensive materials (catalysts like platinum)
• ❌ Hydrogen is difficult and expensive to store
• ❌ Highly flammable and explosive under pressure
• ❌ Few hydrogen filling stations available
• ❌ Less efficient at low temperatures