ironFechemical element, metal of Group VIII VIIIb of the periodic table, the most used and cheapest metal.
Occurrence, uses, and properties.

Iron makes up 5 percent of the Earth’s crust and is second in abundance to aluminum among the metals and fourth in abundance behind oxygen, silicon, and aluminum among the elements. Iron, which is the chief constituent of the Earth’s core, is the most abundant element in the Earth as a whole (about 35 percent) and is relatively plentiful in the Sun and other stars. In the crust the free metal is rare, occurring as terrestrial iron (alloyed with 2–3 percent nickel) in basaltic rocks in Greenland and carbonaceous sediments in the United States (Missouri) and as a low-nickel meteoric iron (5–7 percent nickel), kamacite. Nickel-iron, a native alloy, occurs in terrestrial deposits (21–64 percent iron, 77–34 percent nickel) and in meteorites as taenite (62–75 percent iron, 37–24 percent nickel). (For mineralogical properties of native iron and nickel-iron, see native elements [table].) Meteorites are classified as iron, iron-stone, or stony according to the relative proportion of their iron and silicate-mineral content. Iron is also found combined with other elements in hundreds of minerals; of greatest importance as iron ore are hematite, magnetite, limonite, and siderite. The metal is extracted by smelting with carbon (coke) and limestone. For specific information on the mining and production of iron, see Industries, Extraction and Processing: Iron.


Iron, as commonly available, nearly always contains small amounts of carbon, which are picked up from the coke during smelting. These modify its properties, from hard and brittle cast irons containing up to 4 percent carbon to more malleable low-carbon steels containing less than 0.1 percent carbon. Three true allotropes of iron in its pure form occur. Delta iron, characterized by a body-centred-cubic crystal structure, is stable above a temperature of 1,394° C (2,541° F). Below this temperature there is a transition to gamma iron, which has a face-centred-cubic structure and is paramagnetic (weakly affected by a magnet); its ability to form solid solutions with carbon is important in steelmaking. At 912° C (1,674° F) there is a transition to paramagnetic alpha iron, which is also body-centred cubic in structure. Below 770° C (1,420° F), alpha iron becomes ferromagnetic, indicating a change in electronic structure but no change in crystal structure; above 770° C (its Curie point) it loses its ferromagnetism altogether. Alpha iron is a soft, ductile, gray-white metal of high tensile strength. Natural iron is a mixture of four stable isotopes: iron-56 (91.66 percent), iron-54 (5.82 percent), iron-57 (2.19 percent), and iron-58 (0.33 percent). Iron-57, which exhibits a marked Mössbauer effect (recoil-free gamma-ray resonance absorption), has been used in studying magnetism and hemoglobin derivatives and for making a very precise nuclear clock.


The most important oxidation states of iron are +2 and +3, though a number of +4 and +6 states are known. Iron(II) compounds Compounds of iron in the +2 state are designated ferrous and contain the pale green Fe2+ ion or complex ions. Iron(III) compounds Compounds of iron in the +3 state are called ferric and contain the Fe3+ ion (which is yellow to orange to brown, depending on the extent of hydrolysis) or complex ions.


The action of sulfuric acid on iron results in the formation of two sulfur compounds: iron(II) sulfate, or ferrous sulfate, FeSO4, which is commonly available as the heptahydrate FeSO4·7H2O; and iron(III) sulfate, or ferric sulfate, Fe2(SO4)3. Ferrous sulfate heptahydrate, known in commerce as green vitriol, or copperas, is obtained as a by-product of industrial processes using iron ores that have been treated with sulfuric acid. It serves as a starting material for the manufacture of various other iron(II) ferrous compounds and as a reducing agent. It is also employed in making inks, fertilizers, and pesticides and for iron electroplating. Ferric sulfate is produced on a large scale by adding sulfuric acid and an oxidizing agent (e.g., nitric acid or hydrogen peroxide) to a hot solution of iron(II) ferrous sulfate. It is used to make iron alums and other iron(III) ferric compounds; as a coagulant in water purification and sewage treatment; and as a mordant (fixative) in textile dyeing and printing.

With chlorine, iron forms another group of industrially important compounds: iron(II) chloride, or ferrous chloride, FeCl2; and iron(III) chloride, or ferric chloride, FeCl3. Ferrous chloride is obtained as yellow-green deliquescent (moisture-absorbing) crystals by passing dry hydrogen chloride gas over red-hot iron. It also can be prepared in hydrated form, FeCl2·4H2O, by dissolving metallic iron in hydrochloric acid. It is used in the dye industry as a mordant and as a reducing agent. Ferric chloride is generally prepared from ferrous chloride through the action of chloride or nitric acid. It is used to make many other iron(III) ferric compounds and as a chlorinating agent for silver, copper, and some organic compounds. Treating a solution of trivalent iron Fe3+ with the complex iron cyanide hexacyanoferrate (II)ion, [Fe(CH6CN6)]4-, yields a deep-blue precipitate called Prussian blue (q.v.). This pigment has a slight reddish tint and is employed in paints, enamels, and lacquers.

A number of iron compounds have been found medically useful. For example, ferrous gluconate, Fe(C6H11O7)2·2H2O, and ferric pyrophosphate, Fe4(P2O7)·xH2O, are among the compounds frequently used to treat anemia. Various ferric salts, which act as coagulants, are applied to wounds to promote healing.

atomic number26atomic weight55.847melting point1,538° C (2,800° F)boiling point3,000° C (5,432° F)specific gravity7.86 (20° C)valence2oxidation states+2, +3, +4, +6electronic config.2-8-14-2 or (Ar)[Ar]3d64s2