What is Electrolysis? Core Principles Explained

Electrolysis is defined as the decomposition (breaking down) of an ionic compound into its elements when molten or in aqueous solution by the passage of an electric current[citation:1][citation:9]. It's a fundamental concept tested across all major exam boards.

Key Condition: The compound must be ionic and have mobile ions (not in solid state)[citation:1]. This is why solid sodium chloride doesn't conduct electricity, but molten or dissolved NaCl does[citation:9].

Exam Super-Mnemonic: PANIC & OIL RIG

PANIC: Positive is Anode, Negative Is Cathode. Use this to instantly label electrodes in any diagram[citation:1][citation:2].

OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons)[citation:9]. Oxidation happens at the anode, reduction at the cathode.

Essential Definitions (Write These in Your Notes!)

Electrolyte
The ionic compound in a molten or aqueous state that conducts electricity and is decomposed during electrolysis[citation:1][citation:9]. Examples: molten lead(II) bromide, aqueous sodium chloride.
Electrode
A rod (usually metal or graphite) that conducts electricity into or out of the electrolyte[citation:2].
Anode
The positive electrode. Anions (negative ions) move here to be oxidised (lose electrons)[citation:1][citation:8].
Cathode
The negative electrode. Cations (positive ions) move here to be reduced (gain electrons)[citation:1][citation:8].
Inert Electrode
An unreactive electrode (e.g., graphite/platinum) that does not participate in the reaction[citation:8].
Electroplating
The process of depositing a layer of metal onto an object at the cathode using electrolysis[citation:8][citation:9].

Examiner's Tip: In electrolysis, electrons flow from the anode to the cathode through the wires. Ions are the charge carriers within the electrolyte[citation:2]. Never say "electrons flow through the solution"—that's a common mistake that will cost you a mark!

Predicting Products: The Golden Rules

The number one challenge in electrolysis questions is predicting what forms at each electrode. Use this flowchart in your mind for every question.

1. At the CATHODE (Negative Electrode)

  • Metals or hydrogen are formed here[citation:1][citation:9].
  • If the metal is MORE REACTIVE than hydrogen (e.g., Na, K, Ca, Mg, Al, Zn), then HYDROGEN GAS (H₂) is produced from the reduction of water or H⁺ ions[citation:2].
  • If the metal is LESS REACTIVE than hydrogen (e.g., Cu, Ag, Au), then the SOLID METAL is deposited[citation:2].
  • Half-equation example (for Cu²⁺): Cu²⁺(aq) + 2e⁻ → Cu(s)[citation:9]

2. At the ANODE (Positive Electrode)

  • Non-metals (except hydrogen) are formed here[citation:1][citation:9].
  • If the solution contains HALIDE IONS (Cl⁻, Br⁻, I⁻) in high concentration, the HALOGEN is produced (e.g., Cl₂ gas)[citation:5][citation:9].
  • If NO halide ions are present, or they are very dilute, then OXYGEN GAS (O₂) is produced from the oxidation of hydroxide ions (OH⁻)[citation:5][citation:8].
  • Half-equation example (for OH⁻): 4OH⁻(aq) → 2H₂O(l) + O₂(g) + 4e⁻[citation:9]

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Electrolysis of Aqueous Solutions: The Complete Picture

This is where most students get confused. In aqueous solutions, you must consider the ions from BOTH the compound AND water (H⁺ and OH⁻)[citation:3][citation:5]. The products depend on a competition based on reactivity, concentration, and electrode material[citation:5].

Electrolyte (Aqueous) Cathode Product (Negative) Anode Product (Positive) - Inert Electrodes Key Reason / Observation
Concentrated Sodium Chloride (NaCl) Hydrogen Gas (H₂) Chlorine Gas (Cl₂) High [Cl⁻] favours Cl₂ over O₂[citation:5]. Used industrially to make NaOH, H₂, Cl₂.
Dilute Sodium Chloride Hydrogen Gas (H₂) Oxygen Gas (O₂) Low [Cl⁻] means OH⁻ is oxidised instead[citation:9].
Copper(II) Sulfate (CuSO₄) with inert electrodes Copper Metal (Cu) Oxygen Gas (O₂) Cu²⁺ is less reactive than H⁺, so it deposits. SO₄²⁻ isn't discharged, so OH⁻ is[citation:3][citation:5]. Solution becomes acidic.
Copper(II) Sulfate with COPPER electrodes Copper Metal (Cu) Copper ions (Cu²⁺) Active electrodes: The anode dissolves! Cu(s) → Cu²⁺(aq) + 2e⁻. Used for electroplating and purification[citation:3][citation:5].
Dilute Sulfuric Acid (H₂SO₄) Hydrogen Gas (H₂) Oxygen Gas (O₂) Effectively the electrolysis of water. 2H₂O(l) → 2H₂(g) + O₂(g)[citation:5].

Real-World Link: Industrial Electrolysis

The principles you're learning power the world. The Chlor-Alkali Industry uses concentrated brine electrolysis to make chlorine (for plastics, disinfectants), hydrogen (for fuels), and sodium hydroxide (for soap, paper)[citation:10]. The Hall-Héroult process uses molten Al₂O₃ electrolysis to extract aluminum[citation:10]. Even the purification of copper for electrical wires (99.99% pure) relies on electrolysis with copper electrodes[citation:3].

Frequently Asked Exam Questions (FAQs)

Q1: In the electrolysis of aqueous copper(II) sulfate, why does the blue colour sometimes fade, and sometimes not?

A: It depends on the electrodes! With inert electrodes (graphite/platinum), Cu²⁺ ions are removed from solution and deposited as copper metal at the cathode. This decreases the concentration of blue Cu²⁺(aq) ions, so the colour fades[citation:5]. With active copper electrodes, the anode dissolves to form Cu²⁺ ions at the same rate they are deposited at the cathode. The concentration of Cu²⁺ stays constant, so the blue colour does not fade[citation:3][citation:5].

Q2: How can I remember if a metal or hydrogen is produced at the cathode?

A: Link it to the reactivity series. Draw the series in your mind. If the metal ion is below hydrogen (like Cu²⁺, Ag⁺), the metal is produced. If the metal ion is above hydrogen (like Na⁺, K⁺, Ca²⁺), hydrogen is produced instead[citation:2]. A reactive metal like sodium would instantly react with water anyway, so hydrogen gas is the final product.

Q3: What is "selective discharge" and what controls it?

A: Selective discharge is when, from a mixture of ions, only one type of cation and one type of anion are discharged[citation:8]. It is controlled by three factors: 1) The relative position in the electrochemical series (reactivity), 2) The relative concentration of the ions, and 3) The nature of the electrode (inert vs. active)[citation:5][citation:7].

Q4: What's the difference between an electrolytic cell and a simple (voltaic) cell?

A: This is a classic 6-mark question. An electrolytic cell uses electrical energy to drive a non-spontaneous chemical reaction (e.g., decomposing water). It has an external power supply, and the anode is positive[citation:10]. A simple cell converts chemical energy into electrical energy from a spontaneous reaction (e.g., zinc in copper sulfate). It produces a voltage, and the anode is negative (because oxidation happens there spontaneously)[citation:10].

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