Oxidation and reduction (together called redox) are among the most important ideas in chemistry — they explain rusting, respiration, batteries, the extraction of metals and the reactions of the reactivity series. The two processes always happen together: if one substance is oxidised, another must be reduced. CSEC Chemistry asks you to recognise redox using three definitions and to identify oxidising and reducing agents.
Three ways to define oxidation and reduction
You should know all three, because different questions use different definitions.
1. In terms of oxygen
- Oxidation = gain of oxygen. (e.g. magnesium + oxygen → magnesium oxide — the magnesium is oxidised)
- Reduction = loss of oxygen. (e.g. copper oxide + hydrogen → copper + water — the copper oxide is reduced)
2. In terms of hydrogen
- Oxidation = loss of hydrogen.
- Reduction = gain of hydrogen.
3. In terms of electrons (the most general)
- Oxidation Is Loss of electrons.
- Reduction Is Gain of electrons.
The memory aid is OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons). For example, when magnesium reacts:
Mg → Mg²⁺ + 2e⁻ (magnesium loses electrons — oxidised) Cu²⁺ + 2e⁻ → Cu (copper ions gain electrons — reduced)
Oxidising and reducing agents
- An oxidising agent oxidises another substance — so it gains the electrons and is itself reduced.
- A reducing agent reduces another substance — so it gives away electrons and is itself oxidised.
This wording trips many candidates: the agent does the opposite to its name. The oxidising agent is the one being reduced.
Common oxidising agents include oxygen, chlorine, potassium manganate(VII) and concentrated nitric acid. Common reducing agents include hydrogen, carbon, and reactive metals.
Oxidation states (numbers)
For the General Proficiency you should be able to use simple oxidation numbers to track redox:
- uncombined elements have an oxidation number of 0;
- in compounds, oxygen is usually −2 and hydrogen +1;
- the oxidation numbers in a neutral compound add up to 0.
If an element's oxidation number increases, it has been oxidised; if it decreases, it has been reduced. For example, in changing from Fe²⁺ to Fe³⁺ the iron's oxidation number rises from +2 to +3, so it is oxidised.
Spotting redox in familiar reactions
- Rusting: iron + oxygen + water → hydrated iron(III) oxide. Iron is oxidised (gains oxygen / loses electrons).
- Displacement: a more reactive metal displaces a less reactive one, e.g. zinc + copper(II) sulfate → zinc sulfate + copper. Zinc loses electrons (oxidised); copper ions gain electrons (reduced).
- Combustion of fuels and metals — the fuel/metal is oxidised by oxygen.
- Extraction of metals with carbon — the metal oxide is reduced (loses oxygen) while the carbon is oxidised.
Tests that show redox is happening
A useful indicator is acidified potassium manganate(VII), which is purple. When it acts as an oxidising agent it is reduced and decolourises (turns from purple to colourless). A colour change like this is good evidence of a redox reaction and is worth quoting in answers.
Half-equations and combining them
A powerful way to handle redox is to write a half-equation for each part — one for the oxidation and one for the reduction — showing the electrons. Take the displacement reaction between zinc and copper(II) sulfate:
oxidation: Zn → Zn²⁺ + 2e⁻ reduction: Cu²⁺ + 2e⁻ → Cu
Because the electrons lost by zinc are exactly the electrons gained by the copper ions, the two half-equations can be added together (the electrons cancel) to give the overall ionic equation:
Zn + Cu²⁺ → Zn²⁺ + Cu
Always check that the electrons balance — the number lost must equal the number gained. This is the surest way to confirm a reaction really is redox and to avoid the common error of an equation that does not balance for charge.
Redox and the reactivity series
Redox underlies the whole reactivity series of metals. A reactive metal loses its outer electrons easily, so it is easily oxidised; this is exactly what "reactive" means. That is why:
- a more reactive metal displaces a less reactive one from solution (it pushes its electrons onto the less reactive metal's ions);
- very reactive metals (above carbon, e.g. aluminium) must be extracted by electrolysis, because carbon cannot reduce their oxides;
- less reactive metals (below carbon, e.g. iron, zinc) can be extracted by reduction with carbon in a furnace.
Seeing extraction and displacement as redox — electrons moving from the more reactive to the less reactive species — ties several CSEC topics together and is exactly the kind of link that earns top marks.
Common exam mistakes
- Getting the agent backwards — remember the oxidising agent is itself reduced.
- Forgetting that oxidation and reduction always occur together; you cannot have one without the other.
- Using only the "oxygen" definition when the question involves ions — switch to the electron definition (OIL RIG).
- Writing electron half-equations that do not balance for charge.
Worked example — identifying what is oxidised and reduced
A common question gives a word or symbol equation and asks you to identify the oxidation and reduction. Take the reaction of magnesium with hydrochloric acid:
Mg + 2HCl → MgCl₂ + H₂
- The magnesium starts as an uncombined element (oxidation number 0) and ends as Mg²⁺ in MgCl₂ (oxidation number +2). Its number has risen, so magnesium is oxidised (it lost electrons).
- The hydrogen starts as H⁺ in the acid (oxidation number +1) and ends as H₂ gas (oxidation number 0). Its number has fallen, so hydrogen is reduced (it gained electrons).
So the magnesium is the reducing agent (it gave electrons away) and the hydrogen ions are the oxidising agent. Setting your answer out this way — naming the species, stating its oxidation-number change, and concluding oxidised or reduced — earns full marks and avoids vague statements.
Key terms to remember
- Oxidation — gain of oxygen / loss of hydrogen / loss of electrons.
- Reduction — loss of oxygen / gain of hydrogen / gain of electrons.
- Redox — a reaction in which oxidation and reduction happen together.
- OIL RIG — Oxidation Is Loss, Reduction Is Gain (of electrons).
- Oxidising agent — the substance that oxidises another and is itself reduced.
- Reducing agent — the substance that reduces another and is itself oxidised.
- Oxidation number — a value tracking electron control; a rise = oxidised, a fall = reduced.
- Half-equation — an equation showing the electrons lost or gained at one part of a redox reaction.
Quick recap
- Redox = oxidation and reduction happening together.
- Oxidation = gain of O / loss of H / loss of electrons; reduction = loss of O / gain of H / gain of electrons (OIL RIG).
- An oxidising agent is reduced; a reducing agent is oxidised.
- Rising oxidation number = oxidised; falling = reduced.
- Rusting, displacement, combustion and metal extraction are all redox reactions; acidified KMnO₄ decolourising is evidence of redox.