Rust

Rust, a hydrated form of iron (III) oxide, is the substance formed when iron compounds corrode in the presence of water and oxygen.

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 -> OH-

The hydroxide ions react with the iron (II) ions and more dissolved oxygen to form iron oxide. The hydration is variable (with x water molecules surrounding each iron oxide molecule):

 Fe2+ + 2OH- -> Fe2O3.xH2O

Rusting 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, and inhibited by alkalis.

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. 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 metal 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.