Redox · Lesson 4

Cells we build, and cells we don’t

Corrosion is a galvanic cell nobody asked for; batteries and fuel cells are ones we design on purpose. How iron rusts, how to protect it, and how rechargeable batteries and fuel cells work.

When the Statue of Liberty was unveiled in 1886 it was the colour of a new copper coin. Within about thirty years it had turned green, as its thin copper skin slowly reacted with the air and rain to form a protective patina. The copper was never in danger. The trouble was hidden inside.

The skin was held in place by an iron framework, and wherever iron touched copper, with seawater spray in between, the iron corroded rapidly. Gustave Eiffel’s engineers had insulated the two metals with shellac-soaked asbestos, but over a century the insulation broke down. By the 1980s many of the iron bars had rusted badly, and during the restoration for the statue’s centenary they were replaced, one by one, with stainless steel.

Two different metals in contact through an electrolyte are a galvanic cell, whether you meant to build one or not. This lesson covers corrosion, how to stop it, and the cells we build on purpose: batteries and fuel cells.

How iron rusts

Rusting needs both water and oxygen. A drop of water on steel becomes a tiny galvanic cell:

Where oxygen is scarce (under the middle of the drop), iron is oxidised: the anode, Fe(s) → Fe²⁺(aq) + 2e⁻.
The electrons travel through the metal to where oxygen is plentiful (the edge of the drop): the cathode, O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq).
Fe²⁺ and OH⁻ ions meet in the water and form iron(II) hydroxide, Fe²⁺ + 2OH⁻ → Fe(OH)₂, which is oxidised by more oxygen to hydrated iron(III) oxide, Fe₂O₃.xH₂O: rust.

Rust forms where the ions meet, not where the iron dissolves, which is why a pit develops under the middle of the drop with a ring of rust around it. Dissolved salts speed corrosion up, because they make the water a better conductor of ions; acids speed it up too.

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A drop of water on steel

Watch iron(II) ions and hydroxide ions meet to form a ring of rust. Switch to salt water, then galvanise the steel.

Protecting iron

Method	How it works	Example
Barrier	keeps water and oxygen away from the iron	paint, oil, grease, plastic coating, tin plating
Galvanising	a zinc coating is a barrier, and if scratched, zinc (a stronger reductant) is oxidised instead of iron	roofing iron, buckets, fences
Sacrificial anode	a block of a more reactive metal is connected to the iron and corrodes instead	magnesium or zinc blocks on ship hulls, pipelines, hot-water tanks
Impressed current	the iron is connected to the negative terminal of a power supply, so it is forced to be a cathode	bridges, wharves, buried pipes
Alloying	chromium in stainless steel forms a tough, self-repairing oxide layer	cutlery, sinks, the Statue of Liberty’s new frame

A scratched tin coating is worse than no coating: tin is a weaker reductant than iron, so where the two meet through water the iron becomes the anode and corrodes faster. A scratched zinc coating still protects, because zinc is the stronger reductant.

Batteries: primary and secondary cells

A primary cell (such as an alkaline cell) is used once: its reaction cannot practically be reversed.
A secondary cell is rechargeable. During discharge it acts as a galvanic cell. During recharging, an external supply forces the reaction in reverse, so it acts as an electrolytic cell: the electrode that was the anode becomes the cathode, and vice versa.
In a lithium-ion cell, lithium ions shuttle through the electrolyte between a graphite electrode and a metal oxide electrode, while electrons travel through the external circuit.

Fuel cells

A fuel cell is a galvanic cell whose reactants are supplied continuously, so it never runs down. In a hydrogen fuel cell with an acidic electrolyte:

anode: H₂(g) → 2H⁺(aq) + 2e⁻ cathode: O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l) overall: 2H₂(g) + O₂(g) → 2H₂O(l)

With an alkaline electrolyte the half-equations are H₂(g) + 2OH⁻(aq) → 2H₂O(l) + 2e⁻ and O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq); the overall reaction is the same. The E°_cell is 1.23 V. The electrodes are porous and coated with a catalyst, often platinum, to speed up the reactions.

Advantages: the only product is water; fuel cells convert chemical energy to electrical energy more efficiently than burning fuel in an engine; they run as long as fuel is supplied.
Disadvantages: most hydrogen is still made from fossil fuels; hydrogen is hard to store and transport; platinum catalysts are expensive.

Explaining corrosion protection for full marks

Name the method and state whether it is a barrier, sacrificial or impressed-current method.
For sacrificial protection, compare E° values: the protecting metal is the stronger reductant, so it is oxidised instead of iron.
Write the relevant half-equation, such as Zn → Zn²⁺ + 2e⁻, and state that iron becomes the cathode.

Now put it to workFive worked examples, twenty practice questions and ten harder problems on this lesson.

Go to the questions