Cathode

Summary: The negative electrode in electrolysis where cations are attracted and reduced by gaining electrons, producing metals or hydrogen gas.

Tags: igcse chemistry

Created: 2026-07-18

The cathode is the negative electrode in an electrolytic cell — the electrode connected to the negative terminal of the d.c. power supply. During Electrolysis, positively charged ions (cations) are attracted to the cathode because opposite charges attract, which gives rise to the useful mnemonic “Cations go to the Cathode” (both start with “Ca”). At the cathode, cations gain electrons in a process called Reduction, following the general half-equation Mⁿ⁺ + ne⁻ → M. The products formed at the cathode are always either a metal or hydrogen gas: in molten electrolytes the metal is always deposited, while in aqueous solutions the least reactive cation present is discharged, meaning hydrogen is produced instead of reactive metals such as sodium or potassium. Typical observations at the cathode include a layer of metal being deposited (such as pink-brown copper or shiny grey lead) or bubbles of colourless hydrogen gas. The cathode also plays a central role in Electroplating, where the object to be coated is made the cathode so that metal ions from the solution are reduced and deposited onto its surface.

Role of the Cathode in Electrolysis

In an electrolytic cell, a d.c. power supply drives current through an electrolyte via two electrodes:

The power supply pushes electrons towards the cathode, making it electron-rich and negatively charged. Cations in the electrolyte migrate through the liquid to the cathode, where they collect these electrons and are discharged (converted into neutral atoms).

Key points to remember:

• The cathode is always the negative Electrode in electrolysis.
• Only cations (positive ions) move to the cathode.
• The reaction at the cathode is always reduction — a gain of electrons.
• The products at the cathode are always a metal or hydrogen, never a non-metal like oxygen or chlorine.

Mnemonics:

• Cations → Cathode (both start with “Ca”)
• OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons) — reduction happens at the cathode.
• RED CAT: REDuction at the CAThode.

Reduction at the Cathode

Every cathode reaction is a reduction because the cation gains electrons. The general half-equation for a metal cation is:

$M^{n+}+n{e}^{-}\to M$

Examples of cathode half-equations:

Note that the number of electrons gained matches the charge on the ion: a 2+ ion gains 2 electrons, a 3+ ion gains 3 electrons, and so on. See Ionic Equations and Half Equations for how to construct and balance these.

Molten Electrolytes: The Metal Is Always Produced

When a molten (melted) ionic compound is electrolysed, there are only two types of ion present — the metal cation and the non-metal anion. There is no competition, so:

• The metal is always produced at the cathode.
• The non-metal is always produced at the anode.

Classic example — molten lead(II) bromide (PbBr₂):

• At the cathode: $P{b}^{2+}+2e^{-}\to Pb$ (a shiny grey bead of molten lead forms below the cathode)
• At the anode: $2B{r}^{-}\to B{r}_2+2e^{-}$ (brown fumes of bromine gas)

Industrial example — extraction of aluminium from molten aluminium oxide (dissolved in cryolite):

• At the cathode: $A{l}^{3+}+3e^{-}\to Al$

This is how reactive metals such as aluminium, sodium and potassium are extracted industrially, since they are too reactive to be extracted by reduction with carbon.

Aqueous Electrolytes: The Least Reactive Cation Is Discharged

In aqueous solutions, water provides additional H⁺ ions (and OH⁻ ions), so there is competition at the cathode between the metal cation and hydrogen ions. The rule is:

The cation of the least reactive element is discharged at the cathode.

This is decided using the reactivity series:

Worked examples:

• Electrolysis of copper(II) sulfate solution: copper is below hydrogen in the reactivity series, so copper metal is deposited — a pink-brown coating appears on the cathode. $C{u}^{2+}+2e^{-}\to Cu$
• Electrolysis of concentrated sodium chloride solution (brine): sodium is far more reactive than hydrogen, so sodium ions stay in solution and hydrogen gas is produced instead. $2H^{+}+2e^{-}\to H_2$
• Electrolysis of dilute sulfuric acid: hydrogen gas is produced at the cathode (and collected in a 2:1 volume ratio with the oxygen at the anode). $2H^{+}+2e^{-}\to H_2$

Test for hydrogen: the gas collected at the cathode gives a squeaky pop with a lighted splint.

The Cathode in Electroplating

In Electroplating, a thin layer of metal is deposited onto an object to improve its appearance or resist corrosion. The setup is:

• Cathode: the object to be plated (e.g. a steel spoon to be silver-plated)
• Anode: a bar of the plating metal (e.g. pure silver), which dissolves to replenish the solution
• Electrolyte: a solution containing ions of the plating metal (e.g. silver nitrate)

At the cathode, metal ions from the solution are reduced and deposit as a thin, even coating on the object:

$A{g}^{+}+e^{-}\to Ag$

The same principle is used in the purification (refining) of copper: the impure copper is the anode, the pure copper is the cathode, and copper is transferred from anode to cathode via Cu²⁺ ions in copper(II) sulfate solution. The cathode gains mass while the anode loses mass. Impurities in the anode (such as gold, silver, and platinum) fall to the bottom of the cell as anode sludge and are recovered separately.

Observations at the Cathode

Typical exam-relevant observations:

Cathode vs Anode: Quick Comparison

Electron Flow and the External Circuit

It is important to distinguish between ion migration inside the electrolyte and electron flow in the external circuit:

• Inside the electrolyte: ions move — cations move towards the cathode, anions move towards the anode. Electrons do not travel through the electrolyte.
• In the external circuit (wires and power supply): electrons flow from the positive terminal of the power supply to the cathode, and from the anode back to the negative terminal of the power supply. This means electrons flow into the cathode from the circuit, making reduction possible.

A common exam question asks students to describe the direction of electron flow. The correct answer: electrons flow from the power supply to the cathode, and from the anode to the power supply.

Sources

• Cambridge IGCSE Chemistry syllabus (0620/0971), Topic 4: Electrochemistry
• CGP IGCSE Chemistry Revision Guide — Electrolysis
• Cambridge IGCSE Chemistry Coursebook, Harwood & Lodge
• BBC Bitesize — Electrolysis and electrode reactions
• ChemGuide — Electrolysis of melts and solutions

Common Misconceptions

Related pages: Anode | Electrolysis | Electrode | Reduction | Electroplating | Ionic Equations and Half Equations