Rust is a general term for a series of iron oxides, usually red oxides, formed by the reaction of iron with oxygen in the presence of water or air moisture. Several forms of rust are distinguishable visually and by spectroscopy, and form under different circumstances. Rust consists of hydrated iron(III) oxides Fe2O3·nH2O, iron(III) oxide-hydroxide (FeO(OH), Fe(OH)3. Rusting is the common term for corrosion of iron and its alloys, such as steel. Other metals undergo equivalent corrosion, but the resulting oxides are not commonly called rust. Given sufficient time, oxygen, and water, any iron mass eventually converts entirely to rust and disintegrates.
The oxidation of iron metal
When in contact with water and oxygen iron will rust. If salt is present, for example, in salt water, the iron will rust more quickly. This chemical reaction is used in the production of handwarmers Iron metal is relatively unaffected by pure water or by dry oxygen. As with other metals, a tightly adhering oxide coating, a passivation layer, protects the bulk iron from further oxidation. Thus, the conversion of the passivating iron 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(III) species are formed. Unlike iron(II) oxides, iron(III) oxides are not passivating because these materials do not adhere to the bulk metal. As these iron(III) compounds form and flake off from the surface, fresh iron is exposed, and the corrosion process continues until all of the iron(0) is either consumed or all of the oxygen, water, carbon dioxide, or sulfur dioxide in the system are removed or consumed. 
Chemical reactions associated with rusting
The rusting of iron is an electrochemical process that begins with the transfer of electrons from iron to oxygen. The rate of corrosion is affected by water and accelerated by electrolytes, as illustrated by the effects of road salt (calcium chloride) on the corrosion of automobiles. The key reaction is the reduction of oxygen:
- O2 + 4 e- + 2 H2O → 4 OH-
Because it forms hydroxide ions, this process is strongly affected by the presence of acid. Indeed, 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:
- 2 Fe2+ + 0.5 O2 → 2 Fe3+ + O2−
Additionally, the following multistep acid-base reactions affect the course of rust formation:
- Fe2+ + 2 H2O ⇌ Fe(OH)2 + 2 H+
- Fe3+ + 3 H2O ⇌ 2 Fe(OH)3 + 3 H+
as do the following dehydration equilibria:
- Fe(OH)2 ⇌ FeO + H2O
- Fe(OH)3 ⇌ FeO(OH) + H2O
- 2 FeO(OH) ⇌ Fe2O3 + H2O
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 (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 both serve as an electrolyte, and thus accelerate rust formation, or combine with the hydroxides and oxides of iron to precipitate a variety of Ca-Fe-O-OH species.
Rust is permeable to air and water, therefore the interior iron continues to corrode. Rust prevention thus requires coatings that preclude rust formation. Stainless steel forms a passivation layer of chromium(III) oxide. Similar passivation behavior occurs with magnesium, copper, titanium, and zinc.
An important approach to rust prevention entails galvanization, which typically consists of coating zinc by either hot-dip galvanizing or electroplating. Zinc is traditionally used because it is cheap and adheres well to steel. In more corrosive environments (such as salt water) cadmium is preferred. Galvanization often fails at seams, holes, and joints, where the coating is pierced. In these cases the coating provides cathodic protection to metal, where it acts as a galvanic anode rusting in preference. More modern coatings add aluminium to the coating as zinc-alume, aluminium will migrate to cover scratches and thus provide protection for longer. These approaches rely on the aluminium and zinc oxides protecting the once-scratched surface rather than oxidizing as a sacrificial anode.
Several other methods are available to control corrosion and prevent the formation of rust, colloquially termed rustproofing:
- Cathodic protection makes the iron a cathode in a battery formed whenever water contacts the iron and also a sacrificial anode made from something with a more negative electrode potential, commonly zinc or magnesium. The electrode alone does not react in water but only provides electrons that are otherwise provided by the iron.
- Bluing is a technique that can provide limited resistance to rusting for small steel items, such as firearms; for it to be successful, water-displacing oil is rubbed onto the blued steel.
- Rust formation can be controlled with coatings, such as paint, that isolate the iron from the environment. Large structures with enclosed box sections, such as ships and modern automobiles, often have a wax-based product (technically a "slushing oil") injected into these sections. Such treatments also contain rust inhibitors. Covering steel with concrete provides protection to steel by the high pH environment at the steel-concrete interface.
Rust is associated with degradation of iron-based tools and structures. As rust has higher volume than the originating mass of iron, its buildup can also cause failure by forcing apart adjacent parts — a phenomenon known as "rust smacking." Similarly corrosion of concrete-covered steel and iron can cause the concrete to spall, creating structural problems.
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