First Law of Electrolysis
The mass of substance deposited at an electrode is directly proportional to the quantity of electricity (charge) passed.
Q = charge (C), I = current (A), t = time (s)
Second Law of Electrolysis
The masses of different substances deposited by the same charge are proportional to their chemical equivalent masses (M/z).
F = 96485 C mol⁻¹, M = molar mass (g/mol), z = ion charge
For active anodes (Cu, Ag, Ni etc.) — the anode dissolves as electrolysis proceeds.
First Law: mass is directly proportional to charge Q = It.
When several types of ions are present, the ion requiring the least energy to discharge is preferentially discharged. At the cathode, the ion with the highest reduction potential (E°) is selected first. At an inert anode, the anion that is most easily oxidised is selected. Concentration critically affects this order.
| Metal Ion | Reduction Half-Equation | E° (V) | Deposited? |
|---|
Ions with E° below H⁺ (0.00 V) are usually not deposited in aqueous solution — H₂ gas forms instead.
| Anion | Oxidation Equation | Ease | Notes |
|---|
⚠️ Concentration Effect
Halide ions (Cl⁻, Br⁻, I⁻) are preferentially discharged in concentrated solutions. In dilute solutions, OH⁻ from water is preferentially oxidised, producing O₂. This explains why concentrated brine gives Cl₂ but dilute NaCl(aq) gives O₂.Select ions present in your electrolyte to predict which will be preferentially discharged at each electrode.
Electroplating
Deposits a thin layer of metal onto a substrate to improve appearance, corrosion resistance, or conductivity.
Setup
- Cathode (−): Object to be plated
- Anode (+): Pure plating metal (active — dissolves)
- Electrolyte: Solution of the plating metal's salt
Equations (Copper example)
Industrial Uses
- Gold/silver jewellery plating
- Nickel/chrome on car parts
- Zinc galvanizing of steel
- Tin plating of food cans
Anodizing Aluminium
Thickens the natural oxide layer on aluminium, greatly improving corrosion resistance and allowing the surface to be dyed.
Setup
- Anode (+): Aluminium article (article being anodized)
- Cathode (−): Inert lead or carbon
- Electrolyte: Dilute H₂SO₄(aq) ~15–20%
Equations
Key Features
- Oxide layer is porous — can absorb dyes for colour
- Sealed by boiling water to close pores
- Used in aircraft, construction, consumer electronics
Copper Refining (Purification)
Produces very pure copper (99.99%) essential for electrical wiring, by electrolytic refining of impure blister copper.
Equations
Chlor-Alkali Process
Electrolysis of concentrated brine (NaCl aq) produces three commercially vital products: Cl₂, H₂, and NaOH.
Setup
- Anode (+): Inert titanium
- Cathode (−): Inert steel
- Electrolyte: Concentrated NaCl(aq) — brine
Equations
Uses of Products
- Cl₂: PVC, bleach, pesticides
- H₂: Ammonia synthesis, fuel cells
- NaOH: Soap, paper, aluminium extraction
Hall-Héroult: Aluminium Extraction
Aluminium extracted by electrolysis of molten Al₂O₃ dissolved in molten cryolite at ~960°C. Cannot be reduced by carbon — too reactive.
Setup
- Cathode (−): Carbon lining — liquid Al collects here
- Anode (+): Carbon blocks — burn away in O₂, must be replaced
- Electrolyte: Molten Al₂O₃ in cryolite (lowers m.p. to ~960°C)
Equations
Energy Cost
- ~15 kWh per kg of aluminium produced
- Recycling Al uses only 5% of extraction energy
At 25°C, 1 mol dm⁻³ solutions, 1 atm.
| Ion | Reduction Half-Equation | E° (V) |
|---|
| Anion | Oxidation Equation | Ease |
|---|
| Electrolyte | Cathode | Anode |
|---|
Electrolysis is the decomposition of a compound using electrical energy. An electrolyte (ionic compound in molten or aqueous form) conducts electricity via mobile ions. Electrons flow through external wires; ions migrate through the electrolyte to complete the circuit.
Anode
The positive electrode, connected to the + terminal. Oxidation occurs here — anions lose electrons, or the electrode material itself is oxidised if active. Mnemonic: "AN OX" — ANion OXidation.
Cathode
The negative electrode, connected to the − terminal. Reduction occurs here — cations gain electrons and are deposited. Mnemonic: "RED CAT" — REDuction at CAThode.
Electrolyte
Conducts electricity when molten or dissolved in water due to free mobile ions. Pure water and solid ionic compounds do NOT conduct. Examples: CuSO₄(aq), NaCl(l), H₂SO₄(aq).
Electrode Types
Inert electrodes (graphite C, Pt) do not react with the electrolyte — only transfer electrons. Active electrodes (Cu, Zn, Ni) participate — the anode dissolves into the electrolyte, replenishing metal ions.
Faraday's Constant
F = 96 485 C mol⁻¹ — the charge carried by one mole of electrons. Named after Michael Faraday (1791–1867). 1 F is the charge of 6.022 × 10²³ electrons. Central to all quantitative electrolysis calculations.
Ion Migration
Cations (Cu²⁺, H⁺, Na⁺) migrate toward the cathode (negative). Anions (Cl⁻, SO₄²⁻, OH⁻) migrate toward the anode (positive). Both types carry current through the solution.
Preferential Discharge
At cathode: highest E° discharged first. At anode (inert): discharge order is I⁻ > Br⁻ > Cl⁻ > OH⁻ >> SO₄²⁻. Cl⁻ vs OH⁻ depends on concentration: concentrated Cl⁻ gives Cl₂; dilute gives O₂.
Aqueous vs Molten
In aqueous solutions, water provides H⁺ and OH⁻ that compete for discharge. In molten salts, only ions from the ionic compound are present — the metal and non-metal are obtained directly (e.g. Na and Cl₂ from molten NaCl).
Quantitative Electrolysis
m = MIt/(zF): doubling current or time doubles mass deposited (First Law). For the same Q, the substance with higher M/z deposits more mass — Second Law. Example: Cu²⁺ (z=2, M=63.5) vs Ag⁺ (z=1, M=108).