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Rust

Rust is the oxide that is formed by open-air oxidation of iron. The chemical composition of rust is mainly iron(III) oxide (Fe₂O₃), and under wet conditions may include iron(III) oxide-hydroxide (FeO(OH)). Rusting is the common term for corrosion of iron and its alloys, such as steel. Although oxidation of other metals is equivalent, these oxides are not commonly called rust.

As rust has higher volume than the originating mass of iron, its buildup may force adjacent parts apart - a phenomenon known as rust smacking.

Chemistry

The rusting of iron is one of the more widely used examples of corrosion. This electrochemical process requires the presence of water, oxygen and an electrolyte and leads to the formation of hydrated iron oxides.

Pure, solid iron oxidizes in water:

Fe(s) -> Fe²⁺(aq) + 2e^-

These electrons will quickly react with the disassociated hydrogen ions (in H3O⁺(aq) form) and the dissolved oxygen in the water (O₂(aq)):

4e^-(aq) + 4H₃O⁺(aq) + O₂(aq) -> 6H₂O(l)

Therefore, as seen from the above equation, the more acidic the water, the greater will be the rate of corrosion (since the concentration of H₃O⁺(aq) will be greater.) At extremely low pH’s, the hydrogen ions will react with the electrons producing hydrogen gas instead:

2H⁺(aq) + 2e^-(aq) -> H₂(g)

Thus, as seen from the above equations, the pH of the solution (whether it is pure water or water containing electrolytes) rises. This leads to the formation of OH^- ions (in cases where the body of water is significantly large, the pH does not rise as sharply, but this is of no consequence since OH^- ions are always present, even in pure water.) The iron cations then react with the OH^- or even the H⁺ ions and dissolved oxygen to form a variety of compounds, which constitute rust:

Fe²⁺(aq) + 2OH^-(aq) -> Fe(OH)₂(s)

4Fe²⁺(aq) + 4H⁺(aq) + O₂(aq) -> 4Fe³⁺(aq) + 2H₂O(l)

Fe³⁺(aq) + 3OH^-(aq) -> Fe(OH)₃(s)

From the above equations, it is seen that the pH and amount of dissolved oxygen can affect the outcome of the reactions. In water with limited dissolved oxygen Fe₃O₄(s) is formed, which is a black solid and commonly called lodestone:

6Fe³⁺(aq) + O₂(aq) + 12OH^-(aq) -> 2Fe₃O₄(s) + 6H₂O(l)

The porous Fe(OH)₃ rust can slowly disintegrate into a crystallized form, which is the familiar red-brown rust:

2Fe(OH)₃(s) -> Fe₂O₃•H₂0(s) + 2H₂O(l)

Iron oxide (FeO) can also be formed. The presence of other ions, such as calcium or calcium carbonates reacts with the iron hydroxides and iron oxides to form a variety of precipitates. Other metals corrode via similar chemical processes.

Rust prevention

Hydrated rust is permeable to air and water, allowing the metal to continue to corrode - internally - even after a surface layer of rust has formed. Given sufficient hydration, the iron mass can eventually convert entirely to rust and disintegrate. Corrosion of aluminium is different from steel or iron, in that aluminium oxide formed on the surface of aluminum metal forms a protective, corrosion resistant coating, a process known as passivation. Stainless steel similarly resists rusting by forming a passivation layer of chromium(III) oxide. This is also true of magnesium, copper and zinc.

Galvanization consists of coating metal with a thin layer of another such metal. Typically, zinc is applied by either hot-dip galvanizing or electroplating. Zinc is traditionally used because it is cheap, easy to refine and adheres well to steel. In more corrosive environments (such as at sea) cadmium may be used. 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 rely on the aluminium and zinc oxides protecting the once-scratched surface rather than oxidizing as a sacrificial anode.

There are several other methods 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 itself doesn't react in water, but only to provide electrons to prevent the iron rusting.

• 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 must be rubbed onto the blued steel.

• Corrosion control can be done using a coating to isolate the metal from the environment, such as paint. 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. This may contain rust inhibiting chemicals as well as forming a barrier. Covering steel with concrete provides protection to steel by the high pH environment at the steel-concrete interface. However, if concrete covered steel does corrode, the rust formed can cause the concrete to spall and fall apart. This creates structural problems.

To prevent rust corrosion on automobiles, they should be kept cleaned and waxed. The underbody should be sprayed to make sure it is free of dirt and debris that could trap moisture. After a car is washed, it is best to let it sit in the sun for a few hours to let it air dry. In winter, or in salty conditions, cars should be washed more regularly as salt (sodium chloride) can accelerate the rusting process.

See also

• WD-40
• Cosmoline
• Weathering steel

External links

• A Primer on Rust, thorough rust prevention and removal information
• Rust Removal and Prevention Articles

Contents

• Chemistry
• Rust prevention
• See also
• External links