This mini-lesson walks you through the whole of CCEA Unit 2, section 2.2 — Redox, Rusting & Iron: what oxidation and reduction really mean, how to spot oxidising and reducing agents, why iron rusts and how we stop it, and how iron is extracted in the blast furnace.
Make a guess: what do oxidation and reduction really keep track of?
Mission 1What do 'oxidation' and 'reduction' really mean?
Mission 2Why let one metal corrode to save another?
Mission 3Why put limestone in a blast furnace?
Work through each screen, answer the questions as you go (some are wordy, some pick apart equations) and collect ⭐ stars. Screens flagged Higher go beyond the Foundation core. Press Start when you're ready.
Defining redox · oxygen & hydrogen
Oxidation and reduction — the oxygen way
The first way to define these terms is in terms of oxygen (or hydrogen):
Oxidation is the gain of oxygen (or loss of hydrogen).
Reduction is the loss of oxygen (or gain of hydrogen).
A reaction in which both oxidation and reduction happen is called a redox reaction — they always go together, because one substance can only lose oxygen if another gains it.
Watch out: "reduction" does not mean the amount gets smaller. It means a substance has lost oxygen (or, as we'll see next, gained electrons).
Quick check
Oxidised or reduced?
?In the reaction Fe₂O₃ + 3CO → 2Fe + 3CO₂, what happens to the iron oxide (Fe₂O₃)?
Higher · defining redox · electrons
The electron way — OIL RIG
The deeper definition is in terms of electron transfer. Remember it with OIL RIG:
Zn + Cu²⁺ → Zn²⁺ + Cu. Zinc loses electrons (oxidised); copper ions gain them (reduced).
Common slip: people mix up OIL RIG because losing negative electrons feels like "going down". It doesn't — losing electrons is oxidation; gaining electrons is reduction.
Quick check
Following the electrons
?A magnesium atom changes into a magnesium ion: Mg → Mg²⁺ + 2e⁻. Using OIL RIG, the magnesium has been…
Higher · oxidising & reducing agents
Agents: who does what to whom?
In a redox reaction we name the two reactants by the job they do:
An oxidising agent oxidises the other substance — so it gains the electrons and is itself reduced.
A reducing agent reduces the other substance — so it gives away electrons and is itself oxidised.
The classic trap: the oxidising agent is the one that gets reduced. It sounds backwards, but the oxidising agent causes oxidation in something else by taking that thing's electrons.
Quick check
Name the agent
?In 2Mg + O₂ → 2MgO, the magnesium gains oxygen. Which species is the oxidising agent?
Sort it
Oxidation or reduction?
Tap a half-reaction, then tap the box it belongs in. (Use OIL RIG and the oxygen rules.)
🔺 Oxidation
🔻 Reduction
Mission 1 check
What do 'oxidation' and 'reduction' really mean?
?Zinc oxide is heated with carbon: 2ZnO + C → 2Zn + CO₂. A student says the zinc oxide is reduced because there is less of it at the end. What does reduced really mean here?
Mission 2 of 3
Why let one metal corrode to save another?
Make a guess: why bolt a zinc block onto a steel ship?
The rusting of iron
What rusting needs
Rusting is the slow oxidation of iron. For iron to rust, two things must both be present:
oxygen (from the air), and
water (moisture).
Remove either one and the iron will not rust. The classic test uses three test tubes with a nail in each:
Only tube A (water and air) rusts. B has no air; C has no water — proving both are needed.
Why each control works: boiling the water drives out dissolved air, and the oil layer stops more dissolving (B). The drying agent (e.g. anhydrous calcium chloride) keeps the air dry (C). Salty water makes rusting faster.
The rusting reaction
Rust is hydrated iron(III) oxide
When iron rusts it gains oxygen, so the iron is oxidised. The product, rust, is hydrated iron(III) oxide:
iron + oxygen + water → hydrated iron(III) oxide4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃·xH₂O (rust)
Because the iron loses electrons to oxygen, rusting is a redox reaction — the iron is oxidised and the oxygen is reduced.
Good to know: rust's flaky, crumbly oxide does not stick to the metal underneath, so fresh iron keeps being exposed and the rusting continues — unlike aluminium, whose oxide forms a protective layer.
Quick check
Reading the experiment
?In the three-test-tube experiment, why does the nail in boiled water sealed under a layer of oil not rust?
Preventing rust · barrier methods
Keep the air and water out
Most rust prevention works by putting a barrier between the iron and the air/water, so one of the two conditions can never reach the metal:
Painting — cheap, used on cars, bridges and railings.
Oiling / greasing — for moving parts like tools and chains.
Coating with plastic — e.g. on garden chairs and dish racks.
Galvanising — coating with a layer of zinc.
Galvanising is clever: the zinc is a barrier, and if the coating is scratched the zinc still protects the iron by sacrificial protection — coming up next.
Preventing rust · sacrificial protection
Sacrificial protection
Attach a more reactive metal (such as zinc or magnesium) to the iron. Because it is higher in the reactivity series, it gives up its electrons and corrodes instead of the iron — it is "sacrificed" to save the iron:
Zinc or magnesium blocks bolted to a ship's hull or a pipeline corrode first, protecting the iron beneath.
Must be MORE reactive: sacrificial protection only works if the attached metal is above iron in the reactivity series. Tin (below iron) would do the opposite — it makes the iron rust faster once the coating is scratched.
Match it
Match the prevention method
Tap a method on the left, then tap how it works on the right.
Mission 2 check
Why let one metal corrode to save another?
?A galvanised (zinc-coated) steel bucket gets a deep scratch, so bare steel shows. Months later, the steel at the scratch has still not rusted. Why?
Mission 3 of 3
Why put limestone in a blast furnace?
Make a guess: why put limestone in a blast furnace?
Extraction of iron
Iron from its ore
Iron (used, mostly as steel, in bridges and structures because of its strength) is less reactive than carbon, so it is extracted by reduction with carbon in a blast furnace. Three raw materials go in (plus hot air):
Iron ore — haematite, which is iron(III) oxide, Fe₂O₃ (the source of the iron).
The molten slag (calcium silicate) runs to the bottom and floats on the denser molten iron, so the two can be tapped off separately. The slag is used in road-building and cement.
So the limestone's job is to remove acidic impurities from the iron ore as easily-removed slag — it does not become part of the iron.
Quick check
The reducing agent
?In the blast furnace, which substance directly reduces the iron(III) oxide to iron (Fe₂O₃ + 3CO → 2Fe + 3CO₂)?
Calculate
Balance the equation
1Balance: Fe₂O₃ + __ CO → 2Fe + 3CO₂. What number goes in the blank in front of CO?
CO
Hint: there are 3 oxygen atoms in Fe₂O₃ and you need 6 oxygen atoms in 3CO₂. Count the carbons.
Mission 3 check
Why put limestone in a blast furnace?
?Iron ore from a new mine contains a lot of sand (silicon dioxide, SiO₂). Why does the furnace operator add limestone?
Recap
The facts to know
Oxidation: gain of oxygen / loss of hydrogen / loss of electrons (OIL)
Reduction: loss of oxygen / gain of hydrogen / gain of electrons (RIG)
Oxidising agent: gains electrons → itself reduced
Reducing agent: gives electrons → itself oxidised
Rusting needs: oxygen and water → hydrated iron(III) oxide