A battery, a rusting nail and an aluminium can all depend on electrons moving from one substance to another. Learn how to control that movement.
Every phone battery, every shiny chrome tap and every aluminium can has a connection with chemistry and electricity. In all of them, electrons are transferred from one substance to another. Electrochemistry is the study of that transfer: how chemical reactions can produce electricity, and how electricity can force chemical reactions to happen.
A redox reaction is one in which electrons are transferred. The name joins two words: reduction and oxidation. These always happen together.
Consider zinc placed in copper(II) sulfate solution. The zinc slowly dissolves and a reddish-brown layer of copper forms. In ionic form the reaction is:
We can split this into two half-equations, one for each process:
| Half-equation | Process | What happens to the species |
|---|---|---|
| Zn → Zn²⁺ + 2e⁻ | Oxidation | Zinc atoms lose electrons |
| Cu²⁺ + 2e⁻ → Cu | Reduction | Copper ions gain electrons |
The substance that is oxidised is the reducing agent (it gives away electrons and so reduces the other substance); the substance that is reduced is the oxidising agent. Here zinc is the reducing agent and Cu²⁺ is the oxidising agent. This is the same idea as the reactivity series: a more reactive metal loses electrons more easily, so it can displace a less reactive metal from its compound.
Question: Magnesium reacts with copper(II) sulfate: Mg + CuSO₄ → MgSO₄ + Cu. Write the half-equations and name the oxidising agent.
In every electrochemical cell, oxidation takes place at the anode and reduction takes place at the cathode. This rule never changes. What can change is the electrical sign (polarity) of each electrode:
| Type of cell | Anode | Cathode | Energy change |
|---|---|---|---|
| Galvanic (voltaic) cell, e.g. a battery | Negative (−) | Positive (+) | Chemical energy → electrical energy |
| Electrolytic cell, e.g. electrolysis | Positive (+) | Negative (−) | Electrical energy → chemical energy |
Two helpful memory aids: AN OX and RED CAT (ANode OXidation, REDuction at CAThode), and for electrolysis PANIC (Positive Anode, Negative Is Cathode). In any cell, electrons travel through the external wire from the anode to the cathode, while ions carry the current through the electrolyte.
Not every substance can carry an electric current, and the reason depends on what the charge carriers are.
| Material | Conducts? | Why |
|---|---|---|
| Metals and graphite | Yes, in the solid state | Delocalised electrons are free to move |
| Solid ionic compound (e.g. solid NaCl) | No | Ions are held in fixed positions in the lattice |
| Molten or dissolved ionic compound | Yes | Ions are free to move and carry charge |
| Most covalent compounds (e.g. sugar, pure water) | No (or very poor) | No free ions or electrons |
A compound that conducts electricity when molten or dissolved, and is decomposed by the current, is called an electrolyte. Metals conduct without being chemically changed; electrolytes are chemically changed, because ions are discharged (turned into atoms or molecules) at the electrodes.
A galvanic cell uses a spontaneous redox reaction to produce an electric current. The oxidation and reduction are kept in separate places (called half-cells) so electrons must travel through an external wire. A famous example joins a zinc half-cell to a copper half-cell.
Zinc is more reactive than copper, so it loses electrons more easily. The salt bridge (for example, filter paper soaked in a salt solution) completes the circuit by allowing ions to move; without it the flow of electrons would quickly stop because charge would build up in each beaker. Under standard conditions this zinc–copper arrangement gives about 1.1 volts.
A general pattern helps you predict: the greater the difference in reactivity between the two metals, the larger the voltage. The more reactive metal becomes the negative electrode (anode). A simple cell can even be made from two different metals dipped in salty water or lemon juice, although the current is tiny.
A battery is one or more cells connected together. Different batteries use different chemistry, but all rely on the same principle.
| Type | Typical use | Key idea |
|---|---|---|
| Zinc–carbon or alkaline dry cell | Torches, remote controls, clocks | Zinc is oxidised; single-use (primary) cell |
| Lead–acid battery | Starting a car engine | Lead and lead(IV) oxide plates in sulfuric acid; a 12 V car battery has six 2 V cells; rechargeable |
| Lithium-ion battery | Phones, laptops, electric vehicles | Lithium ions move between electrodes; rechargeable; high energy for its mass |
| Hydrogen fuel cell | Some vehicles and power systems | Hydrogen and oxygen react to give water and electricity: 2H₂ + O₂ → 2H₂O |
In a rechargeable battery, the cell reaction is reversed during charging by connecting an external power supply, so for that period the cell behaves like an electrolytic cell. A fuel cell differs from an ordinary battery because the fuel is supplied continuously from outside, so it does not “run flat” as long as fuel keeps coming. Hydrogen fuel cells give off only water at the point of use, though how the hydrogen is produced still matters for the overall environmental impact.
Do not short-circuit, puncture, crush or heat batteries, and never throw them in a fire; lithium-ion batteries can overheat or catch fire. Do not mix old and new batteries, and take used batteries to a proper recycling point. Keep small button cells away from young children.
Electrolysis is the decomposition of an ionic compound, when molten or in aqueous solution, by passing an electric current through it. The compound is the electrolyte. The electrodes are the metal or graphite rods that carry current in and out. Graphite and platinum are commonly used as inert electrodes because they do not take part in the reaction.
When a molten ionic compound is electrolysed, the only ions present are those of the compound, so the products are easy to predict: the metal forms at the cathode and the non-metal forms at the anode.
| Molten electrolyte | At the cathode (−) | At the anode (+) |
|---|---|---|
| Lead(II) bromide, PbBr₂ | Pb²⁺ + 2e⁻ → Pb (liquid lead) | 2Br⁻ → Br₂ + 2e⁻ (brown bromine vapour) |
| Sodium chloride, NaCl | Na⁺ + e⁻ → Na | 2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas) |
Bromine and chlorine are toxic gases. Electrolysis of molten compounds is done in a fume cupboard by a teacher, never at home.
In a solution, water also supplies hydrogen ions (H⁺) and hydroxide ions (OH⁻), so there is a competition at each electrode. Two simple rules at school level decide which ion is discharged:
With inert electrodes these rules predict the results below. The rules are simplified; concentration and the electrode material can change the outcome, which is why syllabuses often also give the products to learn.
| Electrolyte (inert electrodes) | Cathode product | Anode product | What is left in solution |
|---|---|---|---|
| Concentrated sodium chloride, NaCl(aq) | Hydrogen | Chlorine | Sodium hydroxide solution |
| Dilute sulfuric acid, H₂SO₄(aq) | Hydrogen | Oxygen | Acid becomes slightly more concentrated |
| Copper(II) sulfate, CuSO₄(aq) | Copper (pink-brown layer) | Oxygen | Blue colour fades; solution becomes more acidic |
In the electrolysis of dilute sulfuric acid, hydrogen and oxygen are produced in a volume ratio of 2 : 1, which fits the formula of water, H₂O. Electrolysis of concentrated sodium chloride solution is the basis of the chlor-alkali industry, which makes chlorine (used in disinfecting water and making plastics such as PVC), hydrogen and sodium hydroxide.
Question: Predict the products at each graphite electrode when copper(II) sulfate solution is electrolysed.
Answer: copper at the cathode, oxygen at the anode.
If the anode itself is copper, it takes part. The impure copper anode dissolves (Cu → Cu²⁺ + 2e⁻) and pure copper is deposited on the cathode (Cu²⁺ + 2e⁻ → Cu). The blue colour of the solution stays the same because copper ions are added at the anode as fast as they are removed at the cathode. This is how copper for electrical wiring is purified, because impurities reduce conductivity.
Electroplating coats an object with a thin layer of metal using electrolysis. The object to be plated is the cathode; the anode is a piece of the plating metal; and the electrolyte contains ions of the plating metal. Metal ions are reduced at the object, and the anode slowly dissolves to replace them.
Electroplating is used to protect metals from corrosion (chromium plating, zinc coating), to improve appearance (silver or gold plating) and to give surfaces special properties. The thickness depends on the current and the time: more charge passed means more metal deposited.
Aluminium is a reactive metal, higher in the reactivity series than carbon, so it cannot be extracted from its ore by heating with carbon. Instead, aluminium oxide (purified from bauxite) is electrolysed.
Aluminium oxide has a very high melting point, above 2,000 °C, which would waste huge amounts of energy. So it is dissolved in molten cryolite, which lowers the working temperature to about 950 °C. The electrolyte is then electrolysed between graphite (carbon) electrodes.
| Electrode | Half-equation | Notes |
|---|---|---|
| Cathode (−) | Al³⁺ + 3e⁻ → Al | Molten aluminium sinks and is tapped off |
| Anode (+) | 2O²⁻ → O₂ + 4e⁻ | Oxygen reacts with the hot graphite anodes to make carbon dioxide, so the anodes burn away and must be replaced |
The process uses a very large amount of electricity, so aluminium plants are often located near cheap power. Recycling aluminium needs far less energy than extracting it from the ore, which is why recycling cans and foil is worthwhile.
Question: How many moles of electrons are needed to produce 2 mol of aluminium? How does that compare with 2 mol of copper from Cu²⁺?
Rusting is an oxidation process. Iron reacts with oxygen and water to form hydrated iron(III) oxide, which is reddish-brown and flaky. Both oxygen and water are needed, and salt in the water speeds the process, which is why cars rust faster near the sea or where roads are salted.
| Method | How it works |
|---|---|
| Painting, oiling, plastic coating | Barrier that keeps out oxygen and water |
| Tin plating | Barrier only; once scratched, iron rusts |
| Galvanising (zinc coating) | Barrier, and zinc is more reactive than iron so it is oxidised first (sacrificial protection), even if scratched |
| Sacrificial anodes (magnesium or zinc blocks on ships and pipes) | The more reactive metal corrodes instead of the iron |
Sacrificial protection is a neat example of electrochemistry: the more reactive metal acts as the anode of a small galvanic cell, so iron, as the cathode, is protected.
Try each question first, then tap it to reveal the answer.
Answer: Iron, because it loses electrons (Fe → Fe²⁺ + 2e⁻).
Answer: In the solid the ions are fixed in the lattice; when molten the ions are free to move and carry charge.
Answer: Lead at the cathode; bromine at the anode.
Answer: Al³⁺ + 3e⁻ → Al.
Answer: It dissolves aluminium oxide and lowers the working temperature, saving energy.
Answer: The object to be plated is the cathode; the anode is copper.
Answer: Zinc is more reactive than iron, so it is oxidised in preference to the iron (sacrificial protection).
Suggested use: set up a zinc–copper or magnesium–copper cell with a voltmeter, then repeat the electrolysis of copper(II) sulfate with graphite and with copper electrodes, and ask students to explain the difference. Electrolysis of molten compounds and halogen-producing electrolyses must be teacher demonstrations with proper ventilation. Parents can connect the topic to batteries, rust on bicycles and chrome-plated taps. The simplified discharge rules here match common school syllabuses; please follow your own syllabus for the exact content and notation required.
These reputable sources were consulted to verify the facts in this article:
Copyright note: This article is original EDUSAMBAM educational writing. The diagrams are original inline SVG schematics created for this page and are not copied from any source. Links open external websites; the sources were used to check facts and are suggested for further study.
Last updated: 21 October 2026. Written for students up to age 18 and checked against standard school-level chemistry references. This is general education, not a substitute for your teacher, textbook or syllabus.
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