Anode

Summary: The anode is the positive electrode in an electrolytic cell where oxidation takes place. Negative ions (anions) are attracted to the anode and lose electrons to form neutral atoms or molecules. It is a central concept in electrolysis, electroplating, and electrochemical cells.

Tags: chemistry igcse electrochemistry electrolysis electrode oxidation

Created: 2026-07-13

The anode is defined as the positive electrode in an electrolytic cell, where oxidation — the loss of electrons — takes place. During electrolysis, negatively charged ions known as anions are attracted to the anode because opposite charges attract; once there, the anions surrender their extra electrons to the electrode, becoming neutral atoms or molecules in the process. The identity of the product formed at the anode depends on the electrolyte and the nature of the electrode itself: in the electrolysis of a molten ionic compound, the non-metal element (e.g., chlorine from molten sodium chloride) is always produced at the anode, whereas in aqueous solutions, the product is determined by the relative ease with which different anions lose electrons (for example, hydroxide ions are discharged in preference to sulfate ions). The anode may be constructed from an inert material such as graphite or platinum, which does not participate chemically in the reaction, or from a reactive metal such as copper, which can itself be oxidised and dissolve into the electrolyte — a phenomenon exploited in copper refining and Electroplating. Understanding the anode is fundamental to grasping how electrical energy drives non-spontaneous chemical reactions, linking the abstract concept of electron transfer to practical industrial processes.

Etymology and Naming

The word “anode” was coined by the English scientist Michael Faraday in 1834, derived from the Greek anodos, meaning “way up” or “way upwards.” Faraday conceived the anode as the electrode through which conventional current enters the electrolyte from the external circuit — the “way in” for the current. This historical definition can be a source of confusion because modern understanding describes current as the flow of electrons (which move in the opposite direction to conventional current), but the fundamental relationship remains: the anode is always the electrode at which oxidation occurs.

Location and Charge in Different Cells

The charge and role of the anode depend on the type of electrochemical cell.

In Electrolytic Cells

In an electrolytic cell, the anode is connected to the positive terminal of the external power supply (battery or DC source). Electrons are pulled away from the anode by the power supply, creating a positive charge on the anode that attracts anions from the electrolyte. Oxidation of the anions at the anode releases electrons, which flow through the external circuit towards the cathode.

In Galvanic (Voltaic) Cells

In a galvanic cell, the naming convention differs: the anode is the negative electrode because oxidation still occurs there, generating electrons that flow through the external circuit. The key principle is that the anode is defined by function, not by charge — it is always the electrode where oxidation takes place, regardless of whether the cell is electrolytic or galvanic. For the Cambridge IGCSE syllabus, the focus is on the anode in electrolytic cells, where it is the positive electrode.

Anode Reactions: Oxidation at the Electrode

At the anode, anions give up their extra electrons to the electrode surface. The general half-equation is:

X⁻ → X + e⁻

This is always an oxidation process because the anion loses one or more electrons. The specific half-equation depends on the anion being discharged. Common examples include:

Halide ions (Cl⁻, Br⁻, I⁻) are discharged as diatomic molecules (Cl₂, Br₂, I₂), while oxide and hydroxide ions produce oxygen gas.

Electrolysis of Molten Compounds

When a molten (fused) ionic compound is electrolysed, there is typically only one type of anion present, so the product at the anode is straightforward: the non-metal element is formed.

Example: Electrolysis of molten lead(II) bromide (PbBr₂)

• Anode: 2Br⁻ → Br₂ + 2e⁻
• Brown bromine vapour is released at the anode.

Example: Electrolysis of molten sodium chloride (NaCl)

• Anode: 2Cl⁻ → Cl₂ + 2e⁻
• Greenish-yellow chlorine gas is released at the anode.

In both cases, the anode product is the non-metal element corresponding to the anion in the compound.

Electrolysis of Aqueous Solutions

When the electrolyte is an aqueous solution, water molecules are present alongside the dissolved ions, and a small proportion of water molecules dissociate into H⁺ and OH⁻ ions. This introduces competing reactions at the anode. The anion actually discharged is the one that is more easily oxidised (loses electrons more readily).

The Discharge Series for Anions

The ease of discharge (oxidation) at the anode follows this order, from easiest to hardest:

1. I⁻ (iodide) — easiest to discharge
2. Br⁻ (bromide)
3. Cl⁻ (chloride)
4. OH⁻ (hydroxide)
5. SO₄²⁻ (sulfate), NO₃⁻ (nitrate) — generally not discharged from aqueous solution (oxygen from OH⁻ is discharged instead)

This means that in the electrolysis of aqueous sodium chloride (brine), the chloride ions are discharged at the anode to produce chlorine gas (Cl₂), rather than hydroxide ions producing oxygen, because chloride is more easily oxidised. However, in the electrolysis of aqueous copper(II) sulfate (CuSO₄) with inert electrodes, the sulfate ions are not discharged; instead, hydroxide ions from water are oxidised to produce oxygen gas.

Selected Aqueous Electrolysis Examples

Inert vs. Reactive Anodes

Inert Anodes

An inert anode is made from a material that does not react chemically during electrolysis. Common inert anode materials include:

• Graphite (a form of carbon): cheap, conducts electricity, but degrades slowly over time as it reacts with any oxygen produced, forming CO₂.
• Platinum: highly unreactive and long-lasting, but very expensive; used in laboratory settings.

Inert anodes simply provide a surface for the oxidation of anions; the electrode itself is not consumed (except for the gradual erosion of graphite).

Reactive (Active) Anodes

A reactive anode is made from a metal that can be oxidised during electrolysis. The most important example is a copper anode used in the electrolytic refining of copper. Here, the impure copper anode itself dissolves:

Cu(s) → Cu²⁺(aq) + 2e⁻

The copper ions enter the electrolyte, and pure copper is deposited at the cathode. Impurities (such as silver, gold, and platinum) fall to the bottom as “anode sludge” and are recovered.

Reactive anodes are also essential in electroplating, where the anode is made of the metal to be plated (e.g., a silver anode for silver plating). As the anode dissolves, it replenishes the metal ions in the electrolyte that are being deposited onto the object at the cathode.

The Anode in Electroplating

Electroplating is the process of coating one metal with a thin layer of another metal using electrolysis. The setup is:

• Anode: made of the plating metal (e.g., silver, nickel, chromium)
• Cathode: the object to be plated
• Electrolyte: an aqueous solution of a salt of the plating metal (e.g., silver nitrate for silver plating)

At the anode, the plating metal oxidises and dissolves into the electrolyte: M(s) → Mⁿ⁺(aq) + ne⁻. This maintains the concentration of metal ions in the electrolyte as metal ions are reduced at the cathode and deposited onto the object. The process is self-sustaining as long as current flows.

Observing Anode Products in the Laboratory

IGCSE students are expected to be able to identify anode products through observation and chemical tests:

The Anode and Redox

Oxidation at the anode and reduction at the cathode together constitute a redox reaction. Electrons released by oxidation at the anode travel through the external circuit to the cathode, where they are accepted by cations in a reduction process. The anode is therefore the source of electrons in the external circuit; it is where chemical energy (in a galvanic cell) or electrical energy (in an electrolytic cell) is converted into the flow of electrons that drives the reduction at the cathode.

Practical and Industrial Relevance

• Chlor-alkali industry: Electrolysis of brine using a titanium anode produces chlorine gas, used in disinfectants, PVC, and solvents.
• Aluminium extraction: In the Hall-Heroult Process, a carbon anode is consumed in the electrolysis of molten alumina (Al₂O₃) dissolved in cryolite.
• Electrorefining of copper: Impure copper anodes are dissolved and pure copper is deposited, producing copper of >99.99% purity for electrical wiring.
• Electroplating: Decorative and protective coatings (chrome, silver, gold) are applied using a metal anode that replenishes the plating bath.

Sources

• Cambridge IGCSE Chemistry 0620 Syllabus, Section 4.1: Electrolysis
• Harwood, R. & Lodge, I. (2020). Cambridge IGCSE Chemistry Coursebook (5th ed.). Cambridge University Press.
• Gallagher, R. & Ingram, P. (2021). Complete Chemistry for Cambridge IGCSE (3rd ed.). Oxford University Press.
• BBC Bitesize. “Electrolysis.” Available at: https://www.bbc.co.uk/bitesize/topics/z3qkfrd
• ChemGuide. “Electrolysis.” Available at: https://www.chemguide.co.uk/inorganic/electrolysis/menu.html

Common Misconceptions