Corrosion of Metals (Edexcel GCSE Chemistry)

Oxidation of Metals

• Corrosion is the destruction of materials by chemical substances in their environment which act on them over a period of time

• Most metals can corrode in the presence of oxygen to form the corresponding metal oxide

• Corrosion is caused by redox reactions:

  • The metal loses electrons and is oxidised while the oxygen gains electrons and is reduced

• Rusting is the name given specifically to the corrosion of iron in the presence of water and oxygen from the air:

iron + water + oxygen ⟶ hydrated iron(III)oxide

Rusting of Iron

Barrier Methods

• Rust can be prevented by coating iron with barriers that prevent the iron from coming into contact with water and oxygen

• However, if the coatings are washed away or scratched, the iron is once again exposed to water and oxygen and will rust

• Unlike some other metals, once iron begins to rust it will continue to corrode internally as rust is porous and allows both air and water to come into contact with fresh metal underneath any barrier surfaces that have been broken or scratched

• Common barrier methods include: paint, oil, grease and plastic

Barrier Methods for Preventing Corrosion of Metals

Galvanising / Sacrificial protection

• Iron can be prevented from rusting making use of metals higher in reactivity than iron

• Galvanising is a process where the iron to be protected is coated with a layer of zinc

• ZnCO₃ is formed when zinc reacts with oxygen and carbon dioxide in the air and protects the iron by the barrier method

• If the coating is damaged or scratched, the iron is still protected from rusting because zinc preferentially corrodes as it is higher up the reactivity series than iron

• Compared to iron it loses its electrons more readily:

Zn → Zn²⁺ + 2e^-

• The iron stays protected as it accepts the electrons released by zinc, remaining in the reduced state and thus it does not undergo oxidation

• The electrons donated by the zinc react with hydrogen ions in the water producing hydrogen gas:

2H⁺ + 2e^– → H₂ 

• Zinc therefore reacts with oxygen and water and corrodes instead of the iron

Sacrificial corrosion

• Sacrificial corrosion occurs when a more reactive metal is intentionally allowed to corrode

• An example of this occurs with ships' hulls which sometimes have large blocks of magnesium or magnesium alloys attached

• The blocks slowly corrode and provide protection to the hull in the same way the zinc does by pushing electrons onto the iron which prevents it from being oxidised to iron(III) ions

Electroplating

• Electroplating is a process where the surface of one metal is coated with a layer of a different metal

• The metal being used to coat is a less reactive metal than the one it is covering

• The anode is made from the pure metal used to coat

• The cathode is the object to be electroplated

• The electrolyte is an aqueous solution of a soluble salt of the pure metal at the anode

A piece of iron being electroplated with tin. The electrolyte is tin(II) chloride, a water-soluble salt of tin

Uses of electroplating

• Electroplating is done to make metals more resistant to corrosion or damage, e.g. chromium and nickel plating

• When people talk about a 'tin can', the amount of tin is very small (only about 1%). The can is made from steel and has a very thin coat of tin on the interior surface that resists corrosion from the liquids inside

• It is also done to improve the appearance of metals, e.g. silver plating cutlery and jewellery