Rust

For other uses, see Rust (disambiguation).

Colors and porous surface texture of rust

Rust is a general term for describing iron oxides. In colloquial usage, the term is applied to red oxides, formed by the reaction of iron and oxygen in the presence of water or air moisture. There are also other forms of rust, such as the result of the reaction of iron and chloride in an environment deprived of oxygen, such as rebar used in underwater concrete pillars, which generates green rust. Several forms of rust are distinguishable visually and by spectroscopy, and form under different circumstances.^[1] Rust consists of hydrated iron(III) oxides Fe₂O₃·nH₂O and iron(III) oxide-hydroxide FeO(OH)·Fe(OH)₃.

Given sufficient time, oxygen, and water, any iron mass will eventually convert entirely to rust and disintegrate. Surface rust provides no protection to the underlying iron, unlike the formation of patina on copper surfaces. Rusting is the common term for corrosion of iron and its alloys, such as steel. Many other metals undergo equivalent corrosion, but the resulting oxides are not commonly called rust.

Chemical reactions

File:RustyChainEdit1.jpg

Heavy rust on the links of a chain near the Golden Gate Bridge in San Francisco; it was continuously exposed to moisture and salt spray, causing surface breakdown, cracking, and flaking of the metal.

Oxidation of iron metal

When impure (cast) iron is in contact with water, oxygen, or other strong oxidants, or acids, 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 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 hydroxide species are formed. Unlike iron 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.^[2]

Associated reactions

File:Rust screw.jpg

A rusted (and dirt-encrusted) bolt; note the surface pitting and gradual shape-deformation, caused by severe oxidation

File:Rust from bathtub in Kyiv.jpg

Interior rust in old galvanized iron water pipes can result in brown and black water.

The rusting of iron is an electrochemical process that begins with the transfer of electrons from iron to oxygen.^[3] 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:

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:

4 Fe²⁺ + O₂ → 4 Fe³⁺ + 2 O²⁻

In addition, the following multistep acid-base reactions affect the course of rust formation:

Fe²⁺ + 2 H₂O ⇌ Fe(OH)₂ + 2 H⁺

Fe³⁺ + 3 H₂O ⇌ 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

File:PyOx.JPG

Rusted pyrite cubes embedded in a stony matrix

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-xO_x/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⁺, both of which 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.

A chemical rust indicator can be used for testing the presence of Fe²⁺. Fe²⁺ turns the rust indicator from yellow to blue.

Prevention

Because of the widespread use and importance of iron and steel products, the prevention or slowing of rust is the basis of major economic activities in a number of specialized technologies. A brief overview of methods is presented here; for detailed coverage, see the cross-referenced articles.

Rust is permeable to air and water, therefore the interior metallic iron beneath a rust layer continues to corrode. Rust prevention thus requires coatings that preclude rust formation.

Rust-resistant alloys

See also: Stainless steel and Weathering steel

Stainless steel forms a passivation layer of chromium(III) oxide. Similar passivation behavior occurs with magnesium, titanium, zinc, zinc oxides, aluminium, polyaniline, and other electroactive conductive polymers.

Special "weathering steel" alloys such as Cor-Ten rust at a much slower rate than normal, because the rust adheres to the surface of the metal in a protective layer. Designs using this material must include measures that avoid worst-case exposures, since the material still continues to rust slowly even under near-ideal conditions.

Galvanization

Main article: Galvanization

Galvanization consists of an application on the object to be protected of a layer of metallic zinc by either hot-dip galvanizing or electroplating. Zinc is traditionally used because it is cheap, adheres well to steel, and provides cathodic protection to the steel surface in case of damage of the zinc layer. In more corrosive environments (such as salt water), cadmium plating is preferred. Galvanization often fails at seams, holes, and joints where there are gaps in the coating. In these cases, the coating still provides some partial cathodic protection to iron, by acting as a galvanic anode and corroding itself instead of the underlying protected metal. The protective zinc layer is consumed by this action, and thus galvanization provides protection only for a limited period of time.

More modern coatings add aluminium to the coating as zinc-alume; aluminium will migrate to cover scratches and thus provide protection for a longer period. These approaches rely on the aluminium and zinc oxides re-protecting a once-scratched surface, rather than oxidizing as a sacrificial anode as in traditional galvanized coatings. In some cases, such as very aggressive environments or long design life, both zinc and a coating are applied to provide enhanced corrosion protection.

Plating

Cathodic protection

Main article: Cathodic protection

Cathodic protection is a technique used to inhibit corrosion on buried or immersed structures by supplying an electrical charge that suppresses the electro-chemical reaction. If correctly applied, corrosion can be stopped completely. In its simplest form, it is achieved by attaching a sacrificial anode, thereby making the iron or steel the cathode in the cell formed. The sacrificial anode must be made from something with a more negative electrode potential than the iron or steel, commonly zinc, aluminium, or magnesium. The sacrificial anode will eventually corrode away, ceasing its protective action unless it is replaced in a timely manner.

Cathodic protection can also be provided by using a special-purpose electrical device to appropriately induce an electric charge on the metal to be protected.

Coatings and painting

Main article: Rustproofing

File:Love by SillyPuttyEnemies.jpg

Flaking paint, exposing a patch of surface rust on sheet metal

Rust formation can be controlled with coatings, such as paint, lacquer, or varnish 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 usually also contain rust inhibitors. Covering steel with concrete can provide some protection to steel because of the alkaline pH environment at the steel-concrete interface. However rusting of steel in concrete can still be a problem, since expanding rust can fracture or slowly "explode" concrete from within.

As a closely related example, iron bars were used to reinforce stonework of the Parthenon in Athens, Greece, but caused extensive damage by rusting, swelling, and shattering the marble components of the building.

When only temporary protection is needed for storage or transport, a thin layer of oil, grease, or a special mixture such as Cosmoline can be applied to an iron surface. Such treatments are extensively used when "mothballing" a steel ship, automobile, or other equipment for long-term storage.

Special anti-seize lubricant mixtures are available, and are applied to metallic threads and other precision machined surfaces to protect them from rust. These compounds usually contain grease mixed with copper, zinc, or aluminum powder, and other proprietary ingredients.^{[citation needed]}

Bluing

Main article: Bluing (steel)

Bluing is a technique that can provide limited resistance to rusting for small steel items, such as firearms; for it to be successful, a water-displacing oil is rubbed onto the blued steel.

Inhibitors

Main article: Corrosion inhibitor

Corrosion inhibitors, like gas-phase or volatile inhibitors, can be used to prevent corrosion inside sealed systems. They are not effective when air circulation disperses them, and brings in fresh oxygen and moisture.

Humidity control

Economic impact

Main article: Corrosion

File:Silver Bridge collapsed, Ohio side.jpg

The collapsed Silver Bridge, as seen from the Ohio side

Rust is associated with degradation of iron-based tools and structures. As rust has a much higher volume than the originating mass of iron, its build-up can also cause failure by forcing apart adjacent parts — a phenomenon sometimes known as "rust smacking". It was the cause of the collapse of the Mianus river bridge in 1983, when the bearings rusted internally and pushed one corner of the road slab off its support. Rust was also an important factor in the Silver Bridge disaster of 1967 in West Virginia, when a steel suspension bridge collapsed in less than a minute, killing 46 drivers and passengers on the bridge at the time.

File:Collapsed Kinzua Bridge.jpg

The Kinzua Bridge after it collapsed

The Kinzua Bridge in Pennsylvania was blown down by a tornado in 2003, largely because the central base bolts holding the structure to the ground had rusted away, leaving the bridge anchored by gravity alone.

Like exposed steel, reinforced concrete is also vulnerable to rust damage. Internal pressure caused by expanding corrosion of concrete-covered steel and iron can cause the concrete to spall, creating severe structural problems. It is one of the most common failure modes of reinforced concrete bridges and buildings.

Cultural symbolism

Rust is a commonly-used metaphor for slow decay, since it gradually converts robust iron and steel metal into a soft crumbling powder. A wide section of the industrialized American Midwest and American Northeast, once dominated by steel foundries, the automotive industry, and other manufacturers, has experienced harsh economic cutbacks that have caused the region to be dubbed the "Rust Belt".

In music, literature, and art, rust is associated with images of faded glory, neglect, decay, and ruin.

References

↑ "Interview, David Des Marais". http://nasa.gov/centers/ames/multimedia/audio/MER/mer13.html.
↑ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
↑ Script error edit

External links

40x40px

Wikimedia Commons has media related to: Rust