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Oxidation and Reduction Reactions Explained

Redox reactions power the batteries in your devices, cause iron to rust, allow your cells to extract energy from food, and drive industrial processes from aluminium smelting to the chlorination of drinking water. Understanding how electrons transfer between substances is one of the most practical skills in chemistry.

The Original Definition: Oxygen Transfer

Historically, oxidation meant combining with oxygen, and reduction meant losing oxygen. When iron rusts — 4Fe + 3O₂ → 2Fe₂O₃ — the iron is oxidised. When iron oxide is reduced in a blast furnace using coke (carbon) — Fe₂O₃ + 3CO → 2Fe + 3CO₂ — the iron is reduced back to metal. These reactions always occurred in pairs, which led to the term redox (reduction-oxidation).

The Modern Definition: Electron Transfer

The broader and more useful definition is in terms of electrons. Oxidation is the loss of electrons; reduction is the gain of electrons. The memory aid OIL RIG (Oxidation Is Loss, Reduction Is Gain) is universally taught because it is easy to confuse the two.

In every redox reaction, the two processes must occur simultaneously: you cannot have one substance lose electrons without another gaining them. The substance that loses electrons is oxidised; the substance that gains them is reduced.

• The substance that causes another to be oxidised (by accepting that substance’s electrons) is called the oxidising agent. It is itself reduced in the process.
• The substance that causes another to be reduced (by donating electrons) is called the reducing agent. It is itself oxidised in the process.

Oxidation States

Oxidation states (also called oxidation numbers) are a bookkeeping system for tracking electron distribution in compounds. They are assigned using a set of rules:

• Pure elements always have an oxidation state of 0 (e.g., Fe, O₂, Cl₂).
• Monoatomic ions have an oxidation state equal to their charge (e.g., Na⁺ = +1, Fe³⁺ = +3, Cl⁻ = −1).
• Oxygen is almost always −2 in compounds (except in peroxides, where it is −1).
• Hydrogen is almost always +1 in compounds (except in metal hydrides, where it is −1).
• In a neutral compound, all oxidation states sum to zero. In a polyatomic ion, they sum to the ion’s charge.

In a redox reaction, oxidation is an increase in oxidation state (electrons lost → charge becomes more positive); reduction is a decrease in oxidation state (electrons gained → charge becomes more negative).

Half-Equations

Complex redox reactions are easier to analyse by splitting them into two half-equations: one showing the oxidation, the other showing the reduction. Each half-equation must balance for both atoms and charge.

For the reaction between iron(II) ions and manganate(VII) ions in acidic solution:

• Oxidation half: Fe²⁺ → Fe³⁺ + e⁻
• Reduction half: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O

To combine them into a full ionic equation, multiply the oxidation half by 5 (so the 5 electrons cancel): 5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O.

Electrochemical Cells and Standard Electrode Potentials

If the oxidation and reduction half-reactions are physically separated but connected by a wire and a salt bridge, electrons flow through the wire as electric current. This is the basis of a galvanic (voltaic) cell — a battery.

Each half-reaction has a characteristic tendency to accept electrons, measured as a standard electrode potential (E°) in volts, determined under standard conditions (25°C, 1 mol/L, 1 atm). The half-reaction more likely to gain electrons (higher positive E°) becomes the cathode (reduction); the other becomes the anode (oxidation). The cell’s voltage is E°_cathode − E°_anode. A positive cell voltage means the reaction is spontaneous.

Electrolysis: Forcing Non-Spontaneous Redox

By applying an external electric current, non-spontaneous redox reactions can be driven. This is electrolysis. In the electrolysis of water, an electric current forces water molecules to decompose: water is oxidised at the anode (producing O₂) and hydrogen ions are reduced at the cathode (producing H₂). Industrial electrolysis is used to produce aluminium metal from molten aluminium oxide, to purify copper, and to generate chlorine and sodium hydroxide from salt water (the chlor-alkali process).

Redox in Biology

Living cells are powered by redox chemistry. In cellular respiration, glucose is progressively oxidised — its carbon atoms lose electrons step by step through glycolysis and the Krebs cycle — and the electrons are ultimately transferred to oxygen, which is reduced to water. The energy released at each electron transfer step is captured to synthesise ATP. In photosynthesis, the reverse happens: water is oxidised (releasing O₂), and carbon dioxide is reduced to build glucose using light energy. Vitamins C and E act as antioxidants in the body by donating electrons to neutralise reactive molecules (free radicals) that would otherwise oxidise and damage proteins and DNA.

Summary

Oxidation is the loss of electrons (and an increase in oxidation state); reduction is the gain of electrons (and a decrease in oxidation state). They always occur together — hence redox. The oxidising agent accepts electrons and is itself reduced; the reducing agent donates electrons and is itself oxidised. Oxidation states provide a systematic way to identify which atoms are oxidised or reduced. Half-equations split a redox reaction into its two electron-transfer components. Galvanic cells harvest spontaneous redox reactions to generate electricity; electrolysis uses electrical energy to drive non-spontaneous redox. Redox underpins energy metabolism, photosynthesis, industrial chemistry, and electrochemistry.