Electrochemistry – CIE 0620 Summary
1. What is Electrochemistry?
Electrochemistry is the study of electricity and chemical reactions.
It involves:
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Electrolysis (using electricity to cause a chemical change)
Redox reactions (reduction and oxidation)
2. Conductors and Electrolytes
Electrolyte
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A liquid or molten substance that allows electricity to pass.
Contains free-moving ions.
Examples: molten lead bromide, aqueous sodium chloride.
Conductor
Allows electricity to pass.
Type How electricity flows
Metal
By electrons
Electrolyte By ions
3. Electrolysis
Electrolysis
The breaking down of a compound using electricity.
Requirements for Electrolysis
1. An electrolyte
2. A DC power supply
3. Two electrodes
4. Electrodes
Electrodes
Solid conductors through which electricity enters or leaves the
electrolyte.
Electrode Charge What happens
Positive (+) Oxidation
Anode
Cathode Negative (–) Reduction
👉 Remember:
OIL RIG
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Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)
5. Ions
Ion
A charged particle.
Ion type Charge Example
Cation Positive (+) Na⁺, Cu²⁺
Anion Negative (–) Cl⁻, OH⁻
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Cations move to the cathode
Anions move to the anode
6. Electrolysis of Molten Lead(II) Bromide
Electrolyte
Molten PbBr₂
Ions present
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Pb²⁺
Br⁻
At the Cathode (–)
Pb²⁺ + 2e⁻ → Pb
➡ Lead metal is formed (reduction)
At the Anode (+)
2Br⁻ → Br₂ + 2e⁻
➡ Bromine gas is released (oxidation)
7. Electrolysis of Aqueous Solutions
In aqueous solutions, water also provides ions:
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H⁺
OH⁻
So, different ions compete for discharge.
8. Rules for Discharge of Ions
At the Cathode (–)

Hydrogen is produced unless the metal is less reactive than
hydrogen (e.g. copper, silver).
Example:

Cu²⁺ is discharged instead of H⁺.
At the Anode (+)
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Halide ions (Cl⁻, Br⁻, I⁻) are discharged.
Otherwise, oxygen is produced from OH⁻.
9. Electrolysis of Aqueous Sodium Chloride
(Brine)
Ions present
Na⁺, Cl⁻, H⁺, OH⁻
Cathode
2H⁺ + 2e⁻ → H₂
➡ Hydrogen gas formed
Anode
2Cl⁻ → Cl₂ + 2e⁻
➡ Chlorine gas formed
Solution left
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Sodium hydroxide (NaOH)
📌 Industrial importance:
Used to make:
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Chlorine
Hydrogen

Sodium hydroxide
10. Electrolysis of Dilute Sulfuric Acid
Cathode
Hydrogen gas
Anode
Oxygen gas
Volume ratio
Hydrogen : Oxygen = 2 : 1
11. Half-Equations
Half-equations show electron transfer.
Example:
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Reduction: Cu²⁺ + 2e⁻ → Cu
Oxidation: 2Cl⁻ → Cl₂ + 2e⁻
12. Redox Reactions
Oxidation
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Loss of electrons
Gain of oxygen
Loss of hydrogen
Reduction
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Gain of electrons
Loss of oxygen
Gain of hydrogen
13. Uses of Electrolysis
1. Extraction of metals (aluminium)
2. Electroplating
3. Purification of metals
4. Manufacture of chemicals
14. Electroplating
Electroplating
Using electricity to coat an object with a metal.
Setup
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Object to be plated → Cathode
Plating metal → Anode
Electrolyte → Solution of metal ions
Purpose
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Prevent corrosion
Improve appearance
15. Electrolysis and Reactivity Series
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More reactive metals cannot be extracted by electrolysis of
aqueous solutions.
Less reactive metals are discharged more easily.