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Rust is a general term for a series of iron oxides, usually red oxides, formed by the reaction of iron and oxygen in the presence of water or air moisture. 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). 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.
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When in contact with water and oxygen, or other strong oxidants and/or acids, iron will rust. If salt is present as, for example, in salt water, it tends to rust more quickly, as a result of the electro-chemical reactions. 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. [2]
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:
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:
The following redox reaction also occurs in the presence of water and is crucial to the formation of rust:
Additionally, 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 (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.
Nanoparticles of rust have been shown to be effective at cheaply removing arsenic from water sources to help make them safe to drink. [4]
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, titanium, zinc, aluminium.
An important approach to rust prevention entails galvanization, which typically consists of an application, on the object to be protected, of a layer of zinc by either hot-dip galvanizing or electroplating. Zinc is traditionally used because it is cheap, adheres well to steel and provides a cathodic protection to the steel surface in case of damage of the Zinc layer. 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. In some cases, very aggressive environments or long design life, both zinc and a coating are applied to provide corrosion protection.
Several other methods are available to control corrosion and prevent the formation of rust, colloquially termed rustproofing:
Cathodic protection is a technique used to inhibit corrosion on buried or immersed structures by supplying an electrical charge that suppresses the electrochemical reaction. If correctly applied, corrosion can be stopped completely. In its simplest form it is achieved by attaching a sacrificial anode thus 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.
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.
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.
Another method to avoid rust is to control the environment. Controlling the humidity, if possible, below a certain thereshold can reduce or stop the corrosion process.
Corrosion inhibitors, like gas phase or volatile inhibitors can be used to prevent corrosion in closed systems.
A simple and inexpensive way to remove rust from steel surfaces by hand is to rub the steel with aluminium foil dipped in water. Aluminium has a higher reduction potential than the iron in steel, which may help transfer oxygen atoms from the iron to the aluminium. The aluminium foil is softer than steel and will not scratch it, as steel wool will, but as the aluminium oxidizes, the aluminium oxide produced becomes a fine metal polishing compound. [5]
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.
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 resting by gravity alone.
Similarly 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.
Rust is a city in Burgenland, Austria.
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'RUST (O.E.' rust, a word which appears in many Teutonic languages, cf. Du. roest, Ger. rost); in origin it is allied with "ruddy" and "red," the reddish-brown powdery substance which forms on the surface of iron or steel exposed to atmospheric corrosion. Formerly the process was regarded as oxidation pure and simple, and, although it was known that iron did not rust in dry air, yet no attempt was made to explain why water was necessary to the action. F. Crace-Calvert in 1871 showed that the carbon dioxide of the atmosphere was a factor; and in 1888 Crum Brown published the theory - termed the "carbonic acid theory" - that water and carbon dioxide react with iron to form ferrous carbonate and hydrogen, the ferrous carbonate being subsequently oxidized by moist oxygen to ferric hydrate and regenerating carbon dioxide, which again reacts with more iron. This theory was controverted by Wyndham Dunstan, who attempted to prove that carbon dioxide was not necessary to rusting; and in place of the acid theory, he set up a scheme which involved the production of hydrogen peroxide. G. T. Moody has since shown that when all traces of carbon dioxide are removed (which is a matter of great experimental difficulty) iron may be left in contact with oxygen and water for long periods without rust appearing, but on' the admission of carbon dioxide specks are rapidly formed. It also appears that rust changes in composition on exposure to the atmosphere, both the ferrous oxide and carbonate being in part oxidized to ferric oxide. Acids, other than carbonic, may promote rusting; this is particularly the case with ironwork exposed to the acids - sulphurous, nitric, &c. - contained in smoke. It is probable that the action depends upon the presence of iron, oxygen and water, and some acid which makes the water an electrolyte.
Steel differs in many ways from iron in respect of atmospheric corrosion; the heterogeneous nature of steel gives occasion to a selective rusting, ferrite is much more readily attacked than the cementite and pearlite; moreover, the introduction of other elements may retard rusting; this is particularly the case with the nickel-steels.
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Rust is a type of corrosion. Rust is the slow breaking down of metals, due to air or water coming into contact with them. Almost all metals rust, but some metals are protected by a thin paint. Alloys (mix of metals) such as stainless steel, rust much slower than iron. When a piece of metal rusts, it changes to a different color, (ex. iron turns red/brown) and the metal falls apart. Rust appears on metal if it is left outside in the damp air. For example rust occures mostly in cracks, or where two parts meet on metal. If it rains the water will enter the cracks and make it easier to rust because it would be hard to remove the water from the cracks. So since the water stays in the cracks, the metal starts to corrode. Which at the end it turns into rust. The iron cannot be used or reused as it is completely damaged once it rusts.
Some metals, such as aluminum, form a very very thin coating of corrosion on the metal. The metal cannot continue corroding since the coating blocks the rest of the metal from corroding. That is why aluminum keeps its shine. It also makes aluminum seem very unreactive, even though it is reactive enough to react with water!
)]] Rust can also be any of various basidiomycete fungi that are parasitic on plants and produce reddish or brownish spots on leaves.
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