Gas anodes, coatings and corrosion protection

The previous trail¹ emphasised the importance of being alert to the unexpected and being prepared to bring out the chemistry of that observation. And what could be more unexpected - paradoxical even - than a gas anode? 

Chemistry in the pipeline

The photo (right) was taken on Lepe Beach at Southampton Water. It was close to where the undersea pipeline to the Isle of Wight enters the water. A steel pipeline would corrode in the marine environment. Without protection, iron atoms in the pipe would lose electrons, becoming iron(II) and then iron(III) ions. This could occur under-sea, remote from the air. At this part of the pipeline, the iron is acting as an anode: 

Fe(s) + aq → Fe²⁺(aq) + 2e^- (1) 

Fe²⁺(aq) → Fe³⁺(aq) + e^- (2) 

The electrons would be conducted through the steel to parts of the pipeline well supplied with oxygen, ie near the surface. Here the electrons would be given up to water molecules, which become hydroxide ions. This part of the pipeline is acting as a cathode: 

O₂(aq) + 2H₂O(l) + 4e^- → 4OH^-(aq) (3) 

Note that the part of the pipeline which disintegrates is remote from the air and it would be difficult to spot corrosion taking place. 

To protect the pipeline from corrosion another, more reactive, metal is sacrificed in its place. This metal - a zinc alloy, magnesium or, perhaps surprisingly, aluminium -acts as the 'sacrificial anode'. However, the iron (ie the cathode) is protected. Reactions (4a and 4b) take place rather than reactions (1) and (2). 

Mg(s) + aq → Mg²⁺(aq) + 2e^- (4a) 

Zn(s) + aq → Zn²⁺(aq) + 2e^- (4b) 

Coatings, corrosion and cathodic protection

The zinc coating on steel sheets ('galvanised iron') used for fencing and roofing materials, rubbish bins etc works in exactly the same way. The coating of zinc on a galvanised iron lamp post gives some physical protection but even if the coating is scratched, the more reactive zinc corrodes preferentially, preserving the iron from rusting. In the case of a lamp post, the zinc covers the whole of the iron surface. The zinc is still acting as a sacrificial anode, there is still cathodic protection of the iron but the zinc is physically distributed over the surface of the iron, rather than as a separate 'wire'.  

Zinc is not used to protect steel sheets used for making cans for food because zinc is slightly toxic. Instead, tin is used. Unfortunately, tin is a less reactive metal than iron so that if the coating of tin does get damaged (perhaps as a result of a heavy blow caused by careless handling which results in a major dent to the can) you have ideal conditions for corrosion. The iron dissolves, as in reactions (1) and (2) and the can eventually springs a leak. If the contents are somewhat acidic (eg, canned fruit) instead of reaction (3), reaction (5) may take place, leading to a build up of pressure and possibly a spectacular blowout. 

2H⁺(aq) + 2e^- → H₂(g) (5).