d oxidation state, as part of a reaction. The best-known of these reactions involve oxygen, hence the name "oxidation". The terms "rust" and "rusting" only mean oxidation of iron and its resulting products. Many other oxidation reactions exist which do not involve iron or produce rust. But only iron or alloys that contain iron can rust. However, other metals can corrode in similar ways.
The main catalyst for the rusting process is water. Iron or steel structures might appear to be solid, but water molecules can penetrate the microscopic pits and cracks in any exposed metal. The hydrogen atoms present in water molecules can combine with other elements to form acids, which will eventually cause more metal to be exposed. If chloride ions are present, as is the case with saltwater, the corrosion is likely to occur more quickly. Meanwhile, the oxygen atoms combine with metallic atoms to form the destructive oxide compound. As the atoms combine, they weaken the metal, making the structure brittle and crumbly.
Oxidation of iron
When iron is in contact with water and oxygen it rusts. If salt is present, for example in seawater or salt spray, the iron tends to rust more quickly, as a result of electrochemical reactions. Iron metal is relatively unaffected by pure water or by dry oxygen. As with other metals, like aluminium, a tightly adhering oxide coating, a passivation layer, protects the bulk iron from further oxidation. The conversion of the passivating ferrous oxide layer to rust results from the combined action of two agents, usually oxygen and water.
Other degrading solutions are sulfur dioxide in water and carbon dioxide in water. Under these corrosive conditions, iron hydroxide species are formed. Unlike ferrous oxides, the hydroxides do not adhere to the bulk metal. As they form and flake off from the surface, fresh iron is exposed, and the corrosion process continues until either all of the iron is consumed or all of the oxygen, water, carbon dioxide, or sulfur dioxide in the system are removed or consumed.
When iron rusts, the oxides take up more volume than the original metal; this expansion can generate enormous forces, damaging structures made with iron. See economic effect for more details.
The rusting of iron is an electrochemical process that begins with the transfer of electrons from iron to oxygen. The iron is the reducing agent (gives up electrons) while the oxygen is the oxidising agent (gains electrons). The rate of corrosion is affected by water and accelerated by electrolytes, as illustrated by the effects of road salt on the corrosion of automobiles. The key reaction is the reduction of oxygen:
Because it forms hydroxide ions, this process is strongly affected by the presence of acid. Likewise, the corrosion of most metals by oxygen is accelerated at low pH. Providing the electrons for the above reaction is the oxidation of iron that may be described as follows:
- Fe → Fe2+ + 2 e−
The following redox reaction also occurs in the presence of water and is crucial to the formation of rust:
- 4 Fe2+ + O2 → 4 Fe3+ + 2 O2−
In addition, the following multistep acid–base reactions affect the course of rust formation:
as do the following dehydration equilibria:
From the above equations, it is also seen that the corrosion products are dictated by the availability of water and oxygen. With limited dissolved oxygen, iron(II)-containing materials are favoured, including FeO and black lodestone or magnetite (Fe3O4). High oxygen concentrations favour ferric materials with the nominal formulae Fe(OH)3−xOx⁄2. The nature of rust changes with time, reflecting the slow rates of the reactions of solids.
Furthermore, these complex processes are affected by the presence of other ions, such as Ca2+, which serve as electrolytes which accelerate rust formation, or combine with the hydroxides and oxides of iron to precipitate a variety of Ca, Fe, O, OH species.
The onset of rusting can also be detected in the laboratory with the use of ferroxyl indicator solution. The solution detects both Fe2+ ions and hydroxyl ions. Formation of 