Rust
This article is about the type of corrosion. For the fungus, see rust (fungus). For the person, see Mathias Rust. For the town in Austria, see Rust, Austria.
Rust is the substance formed when iron compounds corrode in the presence of water and oxygen. It is a mixture of iron oxides and hydroxides.
Iron is found naturally in the ore hematite as iron oxide, and purified iron quickly returns to a similar state when exposed to air and water. This corrosion is due to the oxidation of a metal being an energetically favourable process--energy is given off when rust forms. The process of rusting can be summarised as three basic stages: The formation of iron (II) ions from the metal; the formation of hydroxide ions; and their reaction together, with the addition of oxygen, to create rust.
When an iron compound comes in to contact with a drop of water, an electrochemical process starts. On the surface of the metal, iron is oxidised to iron (II):
Fe -> Fe2+ + 2e-
The electrons released travel to the edges of the water droplet, where there is plenty of dissolved oxygen. They reduce the oxygen and water to hydroxide ions:
2e- + 1/2O2 + H2O -> 2OH-
The hydroxide ions react with the iron (II) ions and more dissolved oxygen to form iron oxide. The hydration is variable, however in it's most general form:
Fe2+ + 2OH- -> Fe(OH)2
4Fe(OH)2 + O2 -> 2(Fe2O3.xH2O) + 2H2O
Hence, rust is hydrated iron(III) oxide. Rusting also tends to happen faster at sea. This is due to the higher concentration of sodium chloride ions in the water, making the solution more conductive. Rusting is also accelerated in the presence of acids, but inhibited by alkalis. Rust can often be removed through electrolysis.
Unfortunately rust is unlike aluminium oxide, which forms a protective coating on aluminium to prevent further oxidation. Hydrated iron oxide is permeable to air and water, meaning that the metal continues to corrode after rust has formed. The iron mass eventually converts entirely to rust, and disintegrates. However there exist a number of ways of stopping, or slowing, this process. Galvanising is coating the metal with a thin layer of another metal, such as zinc, which does form a protective oxide. The two most common processes used to achieve this are hot-dip galvanizing and electrogalvanizing.
Also used are sacrificial metals, attached through a conductor to the metal at risk. As the sacrificial anode is chosen to have a higher electrode potential, it is oxidised in preference to the iron. Electrons conduct to the site attacked by oxygen and water, and reduce oxygen to hydroxide irons, like in normal rusting. However because there are no iron (II) ions to react with the hydroxide ions, no rust is formed.
Other techniques include the coating of the metal in an organic polymer or paint. However these are not so powerful--if the surface is scratched the metal is exposed and rust can still form.
On the other hand, a scratch on a galvanised piece of iron will not lead to rust, since zinc has the higher electrode potential. The zinc layer would act as a sacrificial anode as well. The method of using a more reactive metal to corrode in place of iron is also employed to protect ships' hulls, and underground steel pipes. Magnesium plates or blocks may be affixed onto or connected by wire to the iron, thereby supplying the electrons for the reduction of oxygen. The sacrificial corrosion results in the magnesium oxidising instead of the iron hull or pipe, which would be more costly to repair than just the regular replacement of the magnesium blocks.