Oxidation and reduction — the oxygen way

The first way to define these terms is in terms of oxygen:

• Oxidation is the gain of oxygen.
• Reduction is the loss of oxygen.

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).

Oxidised or reduced?

The electron way — OIL RIG

The deeper definition is in terms of electron transfer. Remember it with OIL RIG:

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.

Following the electrons

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 think of it like a generous friend — the oxidising agent causes oxidation in something else by accepting that thing's electrons.

Name the agent

Oxidation or reduction?

Tap a half-reaction, then tap the box it belongs in. (Use OIL RIG and the oxygen rules.)

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:

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.

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:

Because the iron loses electrons to oxygen, rusting is a redox reaction — the iron is oxidised and the oxygen is reduced.

Spec note: 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.

Reading the experiment

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.

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:

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 the prevention method

Tap a method on the left, then tap how it works on the right.

Iron from its ore

Iron is too unreactive to need electrolysis, 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).
• Coke — mostly carbon (fuel and reducing agent).
• Limestone — calcium carbonate, CaCO₃ (to remove acidic impurities).

Hot air is blasted in at the bottom to burn the coke and supply the oxygen and heat for the reactions.

Inside the blast furnace

Three key reactions extract the iron, and a fourth removes the impurities:

Removing the impurities → slag

The ore contains sandy, acidic impurities (mainly silicon dioxide, SiO₂). The limestone deals with them in two steps:

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.

The reducing agent

Balance the equation

The facts to know

You've covered the whole of CCEA Unit 2.2 — Redox, Rusting & Iron. Press Finish to see your score.