Rusting+SB

=4Fe + 3O2 à 2FeO3=

__//The chemical equations /rust formation.//__ The chemical formula / rust is Fe2O3.nH2O. The **overall chemical equation** / the formation of rust is Iron + water + oxygen rust 4 Fe(s) + 6 H2O(l) + 3 O2(g) 4 Fe(OH)3(s) Iron(III) hydroxide, Fe(OH)3 then dehydrates to produce Fe2O3.nH2O(s) or rust [|www.gogle]. /picturs / rust formation.com
 * 1) 2Fe(s) + 2H2O(l) + O2(g) [[image:http://www.chemicalformula.org/sites/default/files/reaction-arrow.png align="absMiddle"]] 2Fe2+(aq) + 4OH-(aq).
 * 2) Fe2+(aq) + 2OH-(aq) [[image:http://www.chemicalformula.org/sites/default/files/reaction-arrow.png align="absMiddle"]] Fe(OH)2(s).
 * 3) Fe(OH)2(s) =O2=> Fe(OH)3(s).
 * 4) Fe(OH)3(s) =dehydrates=> Fe2O3.nH2O(s) or rust

Iron is found naturally in the [|ore] [|hematite] as [|iron oxide], and metallic [|iron] tends to return to a similar state when exposed to air, (hydrogen, oxygen, nitrogen, etc) and water. This [|corrosion] is due to the [|oxidation] reaction when iron metal returns to an energetically favourable state. Energy is given off when rust forms. The process of rusting can be summarised as three basic stages: The formation of iron(II) ions from the metal; the formation of [|hydroxide] ions; and their reaction together, with the addition of [|oxygen], to create rust. [|Iron] is the main component of [|steel] and the corrosion of steel is observed more frequently, since iron is nearly never used without alloying. When steel contacts water, an [|electrochemical] process starts. On the surface of the metal, [|iron] is [|oxidised] to iron(II): Fe -> Fe2+ + 2e- The [|electrons] released travel to the edges of the water droplet, where there is plenty of dissolved oxygen. They [|reduce] the oxygen and water to hydroxide ions: 4e- + O2 + 2H2O -> 4OH- The [|hydroxide] ions react with the iron(II) ions and more dissolved oxygen to form iron oxide. The hydration is variable, however in its most general form: Fe2+ + 2OH- -> Fe(OH)24Fe(OH)2 + O2 -> 2(Fe2O3.//x//H2O) + 2H2O Hence, rust is hydrated [|iron(III) oxide]. Corrosion tends to progress faster in seawater than fresh water due to higher concentration of [|sodium chloride] ions, making the solution more conductive. Rusting is also accelerated in the presence of [|acids], but inhibited by [|alkalis]. Rust can often be removed through [|electrolysis], however the base metal object can not be restored through this method. Corrosion of [|aluminum] is different from [|steel] or [|iron], in that [|aluminum oxide] formed on the surface of aluminum metal forms a protective, corrosion resistant, coating. Hydrated [|iron oxide] is permeable to air and water, meaning that the metal continues to corrode after rust has formed. The iron mass eventually converts entirely to rust, and disintegrates. There are several methods available to control corrosion and prevent the formation of rust. [|Galvanising] consists of coating metal with a thin layer of another metal, such as [|zinc]. The [|electrochemical potential] of zinc is more negative than steel (or iron) and will provide cathodic protection to the underlying steel. Typically, zinc is applied by either [|hot-dip galvanizing] or electrogalvanizing. A nice thing about galvanising is that a scratch on a galvanised piece of [|iron] will not lead to rust at the scratch. The [|zinc] layer acts as a galvanic anode. [|Cathodic protection] is a method to control corrosion and the formation of rust using electrochemical techniques. Corrosion control can be done using a coating to isolate the metal from the environment. 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 will create structural problems Iron is found naturally in the [|ore] [|hematite]as iron oxide, and purified iron quickly returns to a similar state when exposed to air and water. This corrosion is due to the oxidation of a metal being an energetically favourable process--energy is given off when rust forms. The process of rusting can be summarised as three basic stages: The formation of iron (II) ions from the metal; the formation of hydroxide ions; and their reaction together, with the addition of oxygen, to create rust. When an iron compound comes in to contact with a drop of water, an [|electrochemical] process starts. On the surface of the metal, iron is [|oxidised]to iron (II): Fe -> Fe2+ + 2e-
 * Rust** is the substance formed when [|iron] compounds corrode in the presence of [|water] and [|oxygen]. It is a mixture of [|iron oxides] and [|hydroxides]. Rusting is a common term for [|corrosion], and usually corrosion of [|steel].
 * Rust** is the substance formed when [|iron] compounds corrode in the presence of [|water] and [|oxygen]. It is a mixture of iron oxides and hydroxides.

The electrons released travel to the edges of the water droplet, where there is plenty of dissolved oxygen. They [|reduce] the oxygen and water to hydroxide ions: 2e- + 1/2O2 + H2O -> OH-

The hydroxide ions react with the iron (II) ions and more dissolved oxygen to form iron oxide. The hydration is variable (with //x// water molecules surrounding each iron oxide molecule): Fe2+ + 2OH- -> Fe2O3.//x//H2O

Rusting tends to happen faster at sea. This is due to the higher concentration of sodium chloride ions in the water, making the solution more conductive. Rusting is also accelerated in the presence of acids, and inhibited by alkalis. Unfortunately rust is unlike [|aluminium] oxide, which forms a protective coating on [|aluminium] to prevent further oxidation. Hydrated iron oxide is permeable to air and water, meaning that the metal continues to corrode after rust has formed. However there exist a number of ways of stopping, or slowing, this process. //Galvanising// is coating the metal with a thin layer of another metal, such as [|zinc], which //does// form a protective oxide. The two most common processes used to achieve this are [|hot-dip galvanizing]and electrogalvanizing. Also used are //sacrificial// metals, attached through a conductor to the metal at risk. As the sacrificial metal is chosen to have a higher [|electrode potential], it is oxidised in preference to the iron. Electrons conduct to the site attacked by oxygen and water, and reduce oxygen to hydroxide irons, like in normal rusting. However because there are no iron (II) ions to react with the hydroxide ions, no rust is formed. Other techniques include the coating of the metal in an organic polymer or paint. However these are not so powerful--if the surface is scratched the metal is exposed and rust can still form

What is rusting?
//The rusting of iron// Steel is widely used in the manufacture of cars, white goods and the construction industry because it is mucj stronger than iron. The carbon atoms in steel however, greatly decrease the a In the presence of oxygen and water a series of internal galvanic cells or batteries are created. The carbon impurities become the site of reduction. //Reduction half equation:// 4e- + 2H2O(l) + O2(g) ==> 4OH-(aq) The nail is most easily oxidized at points of stress. ie the tip or the head. At these points the crystal lattice is distorted and the iron atoms are easily oxidized. //Oxidation half equation:// 2Fe(s) 2Fe2+(aq) + 4e- The overall or net equation is 2Fe(s) + 2H2O(l) + O2(g) 2Fe2+(aq) + 4OH-(aq) Fe2+(aq) and OH-(aq) ions migrate through the water by diffusion. Refer to the above diagram. When they meet they combine to produce the precipitate, iron(II) hydroxide, Fe(OH)2, which is further oxidized to iron (III) hydroxide, Fe(OH)3, and finally dehydrated to produce rust. The chemistry of the reaction resulting in the formation of rust can be summarized as follows. bility of iron to resist corrosion.
 * Corrosion** is the loss of metallic properties of a metal due to oxidation and is accompanied by the formation of unwanted products. Copper, iron and aluminum metals all corrode over time loosing strength, lustre and electrical conductivity.
 * Rusting** is the corrosion of iron and readily occurs in the alloy steel. The formation of a reddish brown flakes which loosely adheres to the iron is called rust.
 * Steel**is an alloy made of iron and carbon. The carbon atoms in steel greatly increase the strength of the metal. They prevent the iron atoms in the crystal lattice from slipping over one another.

I used gogle/ bing for the information above.