9.5 Corrosion of Metals
Learning Objectives
- State the three conditions that must be present simultaneously for iron to rust.
- Explain how barrier methods prevent corrosion and give real‑world examples.
- Describe sacrificial (cathodic) protection:
- using the Cambridge reactivity series, and
- in terms of electron loss (redox half‑reactions).
- Identify suitable sacrificial anodes for iron and recognise when protection will fail.
1. Conditions Required for Rust Formation
Rust (hydrated iron(III) oxide) forms only when all three of the following are present at the same time:
- Iron or an iron‑containing alloy (e.g., steel)
- Water (or moisture)
- Oxygen from the air
In many practical situations the rate of rusting is greatly increased by the presence of an electrolyte (e.g., dissolved salts in sea water). The overall reaction can be written as:
\[
\text{Fe} + \text{O}_2 + \text{H}_2\text{O} \;\longrightarrow\; \text{Fe}_2\text{O}_3\cdot x\text{H}_2\text{O}\;( \text{rust})
\]
2. Barrier Methods of Protection
Barrier methods stop corrosion by preventing water and/or oxygen from reaching the metal surface.
| Method | How it works | Typical example |
|---|
| Painting |
Forms a continuous, impermeable film that blocks both water and oxygen. |
Painted bridges, car bodies. |
| Greasing / Oiling |
Creates a hydrophobic layer that repels water and reduces oxygen diffusion. |
Oil‑coated pipelines, machinery bearings. |
| Plastic or rubber coating |
Provides complete physical isolation of the metal from the environment. |
Rubber‑lined water tanks, PVC‑covered electrical cables. |
3. Sacrificial (Cathodic) Protection – The Principle
A more reactive metal (the sacrificial anode) is electrically connected to the metal that needs protection (the cathode). The sacrificial metal oxidises preferentially, supplying electrons that keep the protected metal from losing its own electrons.
3.1 Choosing a Sacrificial Anode – Reactivity Series
The Cambridge reactivity series orders metals by their tendency to lose electrons (oxidise). A metal that is **higher (more reactive) in the series** has a more negative standard electrode potential and will act as the anode.
| Cambridge Order (most → least reactive) |
Metal (symbol) |
E° (V) vs Standard Hydrogen Electrode |
| 1 | K (Kalium) | –2.93 |
| 2 | Na (Sodium) | –2.71 |
| 3 | Ca (Calcium) | –2.87 |
| 4 | Mg (Magnesium) | –2.37 |
| 5 | Al (Aluminium) | –1.66 |
| 6 | Zn (Zinc) | –0.76 |
| 7 | Fe (Iron) | –0.44 |
| 8 | H₂ (Hydrogen reference) | 0.00 |
| 9 | Cu (Copper) | +0.34 |
| 10 | Ag (Silver) | +0.80 |
| 11 | Au (Gold) | +1.50 |
For protecting iron, any metal **above iron** in this list can be used. The most common choices are:
- Zinc (Zn) – E° = –0.76 V
- Magnesium (Mg) – E° = –2.37 V (used where a very strong driving force is required, e.g., offshore platforms)
- Aluminium (Al) – E° = –1.66 V (less common for large structures because it forms a protective oxide layer itself)
3.2 Redox (Electron‑Loss) Explanation
When the sacrificial anode is attached to the iron structure, an electrochemical cell is created. The half‑reactions are:
- Anode (sacrificial metal – oxidation)
\[
\text{Zn(s)} \;\rightarrow\; \text{Zn}^{2+}(aq) + 2e^- \quad\text{or}\quad
\text{Mg(s)} \;\rightarrow\; \text{Mg}^{2+}(aq) + 2e^-
\]
- Cathode (reduction on the protected metal surface)
In the presence of dissolved oxygen the dominant reduction is:
\[
\text{O}_2 + 2\text{H}_2\text{O} + 4e^- \;\rightarrow\; 4\text{OH}^-
\]
- Suppressed iron oxidation (what would happen without protection)
\[
\text{Fe(s)} \;\rightarrow\; \text{Fe}^{2+}(aq) + 2e^- \quad\text{(normally)}
\]
Because electrons are supplied from the sacrificial anode, the iron does **not** lose electrons; the above oxidation is effectively prevented.
The flow of electrons is from the anode (Zn or Mg) through the metal structure to the sites where oxygen is reduced. This makes the iron act as the **cathode** of the cell, so it is protected from rusting.
3.3 Simple Electrochemical Cell Diagram (textual description)
Figure 1 (schematic): A steel pipe (Fe) is shown in contact with a zinc coating. The zinc surface is labelled “Anode (Zn → Zn²⁺ + 2e⁻)”. Arrows indicate electrons travelling through the steel to the pipe surface where “Cathode reaction: O₂ + 2H₂O + 4e⁻ → 4OH⁻” occurs. The electrolyte (water containing dissolved O₂ and possibly salt) surrounds the pipe, completing the circuit.
4. Typical Applications
- Galvanised steel – steel coated with a thin layer of zinc (≈ 5 µm). This provides protection for about 50 years in a normal atmospheric environment.
- Marine fittings and ship hulls – magnesium or zinc anodes bolted to the hull; anodes are replaced when they become visibly corroded.
- Underground pipelines – sacrificial zinc anodes are buried alongside the pipe and are often encased in a protective coating.
- Water‑storage tanks – zinc strips welded to the interior of steel tanks.
5. When Sacrificial Protection Fails
- The anode is gradually consumed; when it is completely oxidised the supply of electrons stops.
- At that point the iron (or other protected metal) reverts to being the anode and rusting resumes.
- Regular inspection (visual check, weight loss measurement, or voltage monitoring) is required; once the anode is exhausted it must be replaced or the structure re‑galvanised.
6. Additional Points Required by the Cambridge Syllabus
- Effect of electrolytes: Salt water increases the conductivity of the electrolyte, dramatically accelerating both the oxidation of the sacrificial anode and the reduction of oxygen. This is why marine environments demand more frequent anode replacement.
- Why aluminium often does not need sacrificial protection: Aluminium forms a thin, adherent Al₂O₃ layer that is itself protective, so it is usually left un‑anodised unless a very aggressive environment is involved.
Key Points to Remember
- Rust requires iron, water and oxygen simultaneously; electrolytes speed the process up.
- Barrier methods work by physically blocking water and/or oxygen from the metal surface.
- In sacrificial (cathodic) protection the anode must be **higher (more negative E°)** than the metal to be protected.
- The sacrificial metal oxidises, releasing electrons that travel to the protected metal where they are used for the reduction of oxygen – the protected metal therefore acts as the cathode.
- When the sacrificial anode is exhausted, protection stops and the underlying metal begins to corrode again.