State the conditions required for the rusting of iron and steel to form hydrated iron(III) oxide

Metals – Corrosion of Metals

Objective (AO1)

State the conditions required for the rusting of iron and steel to form hydrated iron(III) oxide, and describe the main methods of preventing it.

What is rust?

When iron or steel is exposed to the atmosphere it can undergo a slow electro‑chemical reaction called rusting. The solid product is a mixture of hydrated iron(III) oxides, conventionally written as Fe₂O₃·nH₂O (where n = 0–3). The colour is the familiar reddish‑brown flaky coating.

Why does steel rust?

Steel is an alloy whose surface consists largely of iron atoms. Therefore the same chemical processes that rust pure iron also act on steel, so the same four environmental conditions are required.

Overall reaction (simplified)

At a macroscopic level the overall process can be written as

In reality the reaction proceeds at many microscopic sites on the metal surface that together form a micro‑galvanic (electro‑chemical) cell. The two half‑reactions are:

The Fe²⁺ ions produced at the anodic sites combine with the OH⁻ ions from the cathodic sites to give Fe(OH)₂, which is further oxidised to Fe(OH)₃ and finally to the hydrated oxide (rust).

Conditions required for rust formation (AO1)

All four of the following must be present simultaneously for rust to develop on iron or steel:

• Iron/steel surface – provides the metal atoms that can be oxidised.
• Oxygen (O₂) – acts as the electron acceptor in the cathodic half‑reaction.
• Water (H₂O) or moisture – creates a thin electrolyte film and supplies hydroxide ions.
• Electrolyte (dissolved salts) – increases the conductivity of the water film, allowing ions to migrate more rapidly.

Preventing rust – barrier methods (AO1)

Barrier methods stop one or more of the four required conditions from being met.

• Painting, oiling, varnishing, polymer coatings – form a waterproof film that blocks water, dissolved salts and oxygen from reaching the metal surface.
• Plating with a non‑reactive metal (e.g., tin, nickel) – provides an impermeable layer that isolates iron/steel from the environment.
• Galvanising with zinc (sacrificial protection – see below) – a special case of barrier protection in which the coating itself participates in the electro‑chemical process.

Sacrificial protection – zinc galvanising (AO1)

Zinc is more reactive than iron, so it oxidises preferentially:

The electrons released by zinc flow to the underlying iron, preventing the iron from losing electrons (i.e., from rusting). The zinc layer itself corrodes (it is the “sacrificial” anode), while the iron remains largely unoxidised even though all four environmental conditions are still present.

Factors influencing the rate of rusting (AO2)

These factors affect how quickly the micro‑galvanic cells operate, i.e., how fast ions can migrate and how rapidly the redox reactions occur.

1. Amount of moisture – a thicker water film gives better ionic conductivity → faster ion migration.
2. Concentration of electrolytes (e.g., NaCl) – dissolved salts increase conductivity dramatically; sea‑air or road‑salt environments accelerate rusting.
3. Temperature – higher temperature raises kinetic energy of particles, increasing reaction rates.
4. Surface condition – scratches, pits or roughness expose fresh metal, enlarging the anodic area.
5. Presence or absence of protective coatings – coatings that block any of the four conditions markedly reduce the rate.

Other examples of metal corrosion (AO1)

• Copper – forms a green patina of basic copper carbonate (Cu₂CO₃(OH)₂) in moist air.
• Aluminium – rapidly forms a thin, adherent layer of aluminium oxide (Al₂O₃) that actually protects the metal underneath.
• Silver – tarnishes to silver sulfide (Ag₂S) when exposed to hydrogen sulfide in the atmosphere.

These examples illustrate that rusting is one specific case of the broader phenomenon of metal corrosion.

Practical investigation (AO3)

Simple experiment – effect of electrolyte concentration

1. Weigh three identical iron nails (record as m₀).
2. Place each nail in a separate beaker: (a) distilled water, (b) 0.5 M NaCl solution, (c) 1 M NaCl solution.
3. Leave the set up at room temperature for one week, keeping the solutions covered to avoid dust.
4. Remove the nails, rinse quickly with distilled water, dry, and re‑weigh (m₁).
5. Calculate the mass gain (Δm = m₁ − m₀) as a measure of rust formed.

Observations typically show increasing rust with higher salt concentration, demonstrating the role of electrolytes in the rate of rusting. Skills assessed: handling equipment, making accurate measurements, recording data, and drawing conclusions.

Summary table

Key points to remember

• Rusting of iron **and** steel requires all four conditions: iron/steel, oxygen, water, and an electrolyte.
• Removing any one condition (by painting, oiling, plating, or galvanising) can dramatically slow or stop rust formation.
• Rust is a mixture of hydrated iron(III) oxides, not a single pure compound.
• Barrier methods prevent the environment from reaching the metal; sacrificial protection allows a more reactive metal (zinc) to corrode in place of iron.
• Understanding the factors that affect the *rate* of rusting helps explain why some environments (coastal, winter roads) are particularly damaging.
• Corrosion is a general term – other metals corrode in different ways, but the same electro‑chemical principles apply.