H12
Corrosion and Restoration
Corrosion is not confined to ships but is a widespread problem. This chapter looks
more closely at the conditions that cause iron and its alloys to rust and the methods that
may be used to slow or stop this process.
Corrosion
Degradation of metal so that it losses its strength and becomes
unable to fulfil its intended purpose.
Rust
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Observations about
rusting
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Rusting of iron
The reddish-coloured, flaky or porous deposit that forms
on exposed iron and steel.
Rust is hydrated iron (III) oxide Fe2O3.xH2O where x can
vary from about 0.5 to 2.
Both oxygen and water are necessary for rust to form
Salt water accelerates rusting
Impure iron rusts more rapidly than pure iron
Iron rusts more rapidly when attached to a less reactive
metal than when it is on its own
Iron rusts less rapidly when it is attached to a more
reactive metal than when it is on its own
Rust occurs more readily when iron is under mechanical
stress.
A galvanic cell is set up when iron rusts.
 Some spot on the iron surface acts as the anode and
oxidation of iron occurs there
Fe
Fe2+ + 2e The cathode is some impurity in the iron and reduction of
oxygen occurs there
O2 + 2H2O + 4e4OH Electrons move from the anode to the cathode through the
iron (the external circuit)
 Fe2+ ions and OH- ions move through the electrolyte
solution on the surface of the iron to complete the circuit
and combine to form iron (II) hydroxide.
Fe2+ + 2OHFe(OH)2(s)
 The Fe(OH)2 is oxidised.
Fe(OH)2(s) + ½ O2
Fe2O3.H2O + H2O
Explaining the
experimental
observations
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Corrosion of
aluminium
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Preventing rust
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Oxygen is needed to be reduced
Water is needed to complete the circuit
Salt water accelerates rusting because it is more
conducting than water
Impurities in the iron act as cathodes
Contact with less active metals means that electrons do
not have to travel as far from the anode to the cathode
More active metals act as the anode and are oxidised not
iron
The orderly crystal structure is disrupted when iron is
under stress making it easier form iron atoms to be
oxidised.
Aluminium is a more reactive metal than iron
Aluminium does not corrode as rapidly as iron because it
forms an impervious layer of aluminium oxide over its
surface. This layer protects the aluminium metal beneath
from further corrosion
Aluminium is a passivating metal.
Sodium chloride can form a porous mixture with the
surface layer of aluminium oxide in coast areas. This
porous layer does not give aluminium the same protection
against corrosion.
Paint the object
Cover the iron with a thin layer of tin
Cover the iron or steel with a vitreous enamel
Passivate the iron with potassium chromate or phosphoric
acid
Galvanise the iron by covering with a layer of zinc
Surface alloys of stainless steel
Coat the surface layer with a polymer
Cathodic protection using a sacrificial anode or an
applied voltage
Which metal will
corrode?
When two metals are in contact the one with the algebraically
lower electrode potential corrodes more rapidly.
Corrosion at depth
Corrosion should be slow because of the low temperatures and
low concentrations of oxygen.
Much of the corrosion of shipwrecks is caused by the action of
anaerobic bacteria that reduce sulfate ions to sulfide
SO42- + 5H2O + 8eHS- + 9OHIron (II) sulfide and Iron (II) hydroxide forms on the surface
of the steel.
Corrosion in acidic
The reduction of oxygen occurs more rapidly in acidic
seawater
conditions than in neutral or basic seawater.
O2 + 4H= + 4e2H2O
Bacteria that feed off the organic material around a shipwreck
provide the acid conditions.
Long-submerged
artefacts
Generally are in poor condition because:
 Metals have been severely corroded
 Objects are encrusted with deposits of calcium carbonate
or coral
 Porous objects have been impregnated with seawater.
Drying the objects causes the formation of salt crystals that
distort the shape, crack the object or react chemically with it.
Seawater is removed from artefacts by leaching.
Crusty deposits are removed physically or by soaking the
object in a dilute acid solution.
Chloride ions can be trapped in the rust around iron objects as
the insoluble Fe(OH)Cl. This chlorine will accelerate further
corrosion of the iron. The chlorine can be removed by
electrolysis.
Objects made from copper and silver as well as iron can be
restored by electrolysis where the object is the cathode and the
metal ions are reduced.