Electrolysis

Electrolysis uses a direct current (DC) to decompose an electrolyte into its elements.

Key Terms

What Happens During Electrolysis?

1. Positive ions (cations, usually metals or hydrogen) move to the cathode (−)
2. Negative ions (anions, usually non-metals) move to the anode (+)
3. At the electrodes, ions gain or lose electrons and become atoms/molecules
4. At the cathode: cations gain electrons (reduction) → metal deposited or hydrogen gas
5. At the anode: anions lose electrons (oxidation) → non-metal released

Memory aid: Positive ions go to the negative electrode. "Opposites attract!" PANIC: Positive Anode, Negative Is Cathode.

When an ionic compound is melted, the ions become free to move. This is the simplest type of electrolysis.

Example: Molten Lead Bromide ($\text{PbBr}_2$)

At the cathode (−): Lead ions gain electrons → lead metal $\text{Pb}^{2+} + 2e^- \rightarrow \text{Pb}$

At the anode (+): Bromide ions lose electrons → bromine gas $2\text{Br}^- \rightarrow \text{Br}_2 + 2e^-$

Observations:

• Silvery liquid (lead) forms at the cathode
• Brown gas (bromine) forms at the anode

Example: Molten Aluminium Oxide ($\text{Al}_2\text{O}_3$)

At cathode: $\text{Al}^{3+} + 3e^- \rightarrow \text{Al}$ (aluminium metal)

At anode: $2\text{O}^{2-} \rightarrow \text{O}_2 + 4e^-$ (oxygen gas)

Aqueous solutions are more complex because water also produces ions:

$\text{H}_2\text{O} \rightleftharpoons \text{H}^+ + \text{OH}^-$

So in an aqueous solution, there are four types of ions present — two from the compound and two from water.

Rules for Products at Each Electrode

At the cathode (−): Metal ions vs hydrogen ions

• If the metal is more reactive than hydrogen → hydrogen gas is produced
• If the metal is less reactive than hydrogen → the metal is deposited

At the anode (+): Non-metal ions vs hydroxide ions

• If the solution contains a halide (Cl⁻, Br⁻, I⁻) → the halogen is produced
• If no halide is present → oxygen is produced (from OH⁻ ions)

4.1 Sodium Chloride Solution (Brine)

Ions present: Na⁺, Cl⁻, H⁺, OH⁻

Cathode: Na is more reactive than H → hydrogen gas $2\text{H}^+ + 2e^- \rightarrow \text{H}_2$

Anode: Halide present (Cl⁻) → chlorine gas $2\text{Cl}^- \rightarrow \text{Cl}_2 + 2e^-$

In solution: Na⁺ and OH⁻ remain → sodium hydroxide solution

This is incredibly important industrially — from brine, we get three useful products: H₂, Cl₂, and NaOH.

4.2 Copper Sulfate Solution (with carbon electrodes)

Ions present: Cu²⁺, SO₄²⁻, H⁺, OH⁻

Cathode: Cu is less reactive than H → copper metal deposited $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$

Anode: No halide → oxygen gas $4\text{OH}^- \rightarrow 2\text{H}_2\text{O} + \text{O}_2 + 4e^-$

Observation: The blue solution becomes paler as Cu²⁺ ions are removed. Copper is deposited as a pink/brown layer on the cathode.

4.3 Dilute Sulfuric Acid

Ions present: H⁺, SO₄²⁻, H⁺, OH⁻

Cathode: Hydrogen gas $2\text{H}^+ + 2e^- \rightarrow \text{H}_2$

Anode: No halide → Oxygen gas $4\text{OH}^- \rightarrow 2\text{H}_2\text{O} + \text{O}_2 + 4e^-$

Effectively, this is the electrolysis of water. The volume of H₂ is double the volume of O₂ (2:1 ratio from $2\text{H}_2\text{O} \rightarrow 2\text{H}_2 + \text{O}_2$).

Aluminium is too reactive to be extracted by carbon reduction, so electrolysis is used.

The Process

1. Aluminium oxide ($\text{Al}_2\text{O}_3$) is dissolved in cryolite (Na₃AlF₆) to lower the melting point from ~2050°C to ~950°C — this saves energy and cost
2. Carbon anodes and a carbon-lined steel cathode are used
3. DC electricity is passed through the molten mixture

At cathode: $\text{Al}^{3+} + 3e^- \rightarrow \text{Al}$ (molten aluminium sinks to bottom)

At anode: $2\text{O}^{2-} \rightarrow \text{O}_2 + 4e^-$

The oxygen produced at the anode reacts with the carbon electrodes: $\text{C} + \text{O}_2 \rightarrow \text{CO}_2$

This means the carbon anodes gradually burn away and must be regularly replaced.

Why Is Aluminium Expensive?

• High energy costs (large amounts of electricity)
• Replacement of carbon anodes
• Mining and purifying aluminium ore (bauxite)

Method

1. Set up two carbon (graphite) electrodes in a beaker of solution
2. Connect to a DC power supply
3. Observe products at each electrode
4. Test gases produced (splint test for H₂, damp litmus for Cl₂, glowing splint for O₂)

Expected Results

• Remember: reduction at cathode (gain electrons), oxidation at anode (lose electrons) — OIL RIG
• For aqueous solutions, always consider the four ions present
• Half equations must balance for atoms AND charge
• Learn the products for common solutions: NaCl(aq), CuSO₄(aq), H₂SO₄(aq)
• Know why aluminium extraction is expensive (energy + anode replacement)

• Electrolysis decomposes ionic compounds using DC electricity
• Cations go to the cathode (−); anions go to the anode (+)
• Molten compounds: straightforward decomposition into elements
• Aqueous solutions: consider metal reactivity (cathode) and halide presence (anode)
• OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons)