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26 manganese ← iron → cobalt - ↑ Fe ↓ Ru Periodic table
General
Name, symbol, number	iron, Fe, 26
Element category	transition metal
Group, period, block	8, 4, d
Appearance	lustrous metallic with a grayish tinge
Standard atomic weight	55.845(2) g·mol−1
Electron configuration	[Ar] 3d6 4s2
Electrons per shell	2, 8, 14, 2 (Image)
Physical properties
Phase	solid
Density (near r.t.)	7.874 g·cm−3
Liquid density at m.p.	6.98 g·cm−3
Melting point	1811 K (1538 °C, 2800 °F)
Boiling point	3134 K (2862 °C, 5182 °F)
Heat of fusion	13.81 kJ·mol−1
Heat of vaporization	340 kJ·mol−1
Specific heat capacity	(25 °C) 25.10 J·mol−1·K−1
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K 1728 1890 2091 2346 2679 3132
Atomic properties
Crystal structure	body-centered cubic a=286.65 pm; face-centered cubic between 1185–1667 K
Oxidation states	6, 5 [1], 4, 3, 2, 1 [2] (amphoteric oxide)
Electronegativity	1.83 (Pauling scale)
Ionization energies (more)	1st: 762.5 kJ·mol−1
	2nd: 1561.9 kJ·mol−1
	3rd: 2957 kJ·mol−1
Atomic radius	140 pm
Atomic radius (calc.)	156 pm
Covalent radius	125 pm
Miscellaneous
Magnetic ordering	ferromagnetic
	1043 K
Electrical resistivity	(20 °C) 96.1 nΩ·m
Thermal conductivity	(300 K) 80.4 W·m−1·K−1
Thermal expansion	(25 °C) 11.8 µm·m−1·K−1
Speed of sound (thin rod)	(r.t.) (electrolytic) 5120 m·s−1
Young's modulus	211 GPa
Shear modulus	82 GPa
Bulk modulus	170 GPa
Poisson ratio	0.29
Mohs hardness	4.0
Vickers hardness	608 MPa
Brinell hardness	490 MPa
CAS registry number	7439-89-6
Most stable isotopes
Main article: Isotopes of iron iso NA half-life DM DE (MeV) DP 54Fe 5.8% >3.1×1022y 2ε capture  ? 54Cr 55Fe syn 2.73 y ε capture 0.231 55Mn 56Fe 91.72% 56Fe is stable with 30 neutrons 57Fe 2.2% 57Fe is stable with 31 neutrons 58Fe 0.28% 58Fe is stable with 32 neutrons 59Fe syn 44.503 d β- 1.565 59Co 60Fe syn 1.5×106 y β- 3.978 60Co
References
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Iron (pronounced /ˈаɪɚn/) is a chemical element with the symbol Fe (Latin: ferrum) and atomic number 26. Iron is a group 8 and period 4 element. Iron and iron alloys (steels) are by far the most common metals and the most common ferromagnetic materials in everyday use. Fresh iron surfaces are lustrous and silvery-grey in color, but oxidise in air to form a red or brown coating of ferrous oxide or rust. Compared with steel which has a hardness of about 140 Brinell, pure iron is soft, about 80 Brinell.^[1][2]

Iron-56 is the second heaviest stable isotope produced by the alpha process in stellar nucleosynthesis, the heaviest being nickel-62; heavier elements require a supernova for their formation. Iron is the most abundant element in the core of red giants, and is the most abundant metal in iron meteorites and in the dense metal cores of planets such as Earth.

Contents 1 Occurrence 2 Characteristics 3 Allotropes 4 Applications 4.1 Iron compounds 5 History 6 Production of iron from iron ore 6.1 Blast furnace 6.2 Further processes 7 Isotopes 8 Iron in organic synthesis 9 Iron in biology 9.1 Nutrition and dietary sources 9.2 Regulation of iron uptake 10 See also 11 Bibliography 12 References 13 External links

Occurrence

Iron is the sixth most abundant element in the Universe, formed as the final act of nucleosynthesis, by silicon fusing in massive stars. While it makes up about 5% of the Earth's crust, the Earth's core is believed to consist largely of an iron-nickel alloy constituting 35% of the mass of the Earth as a whole. Iron is consequently the most abundant element on Earth, but only the fourth most abundant element in the Earth's crust.^[3] Most of the iron in the crust is found combined with oxygen as iron oxide minerals such as hematite and magnetite.

About 1 in 20 meteorites consist of the unique iron-nickel minerals taenite (35–80% iron) and kamacite (90–95% iron). Although rare, iron meteorites are the major form of natural metallic iron on the Earth's surface.

The red color of the surface of Mars is thought to derive from an iron oxide-rich regolith.

Characteristics

Pure iron is a metal, but is rarely found in this form on the surface of the earth, because it oxidizes readily in the presence of oxygen and moisture. In order to obtain metallic iron, oxygen must be removed from naturally occurring ores by chemical reduction – mainly of the iron ore hematite (Fe₂O₃) by carbon at high temperature. The properties of iron can be modified by alloying it with various other metals (and some non-metals, notably carbon and silicon) to form steels.

Nuclei of iron atoms have some of the highest binding energies per nucleon, surpassed only by the nickel isotope ⁶²Ni. The universally most abundant of the highly stable nuclides is, however, ⁵⁶Fe. This is formed by nuclear fusion in stars. Although a further tiny energy gain could be extracted by synthesizing ⁶²Ni, conditions in stars are unsuitable for this process to be favoured. Elemental distribution on Earth greatly favors iron over nickel, and also presumably in supernova element production.^[4]

Iron (as Fe²⁺, ferrous ion) is a necessary trace element used by almost all living organisms. The only exceptions are several organisms that live in iron-poor environments and have evolved to use different elements in their metabolic processes, such as manganese instead of iron for catalysis, or hemocyanin instead of hemoglobin. Iron-containing enzymes, usually containing heme prosthetic groups, participate in catalysis of oxidation reactions in biology, and in transport of a number of soluble gases. See hemoglobin, cytochrome, and catalase.

Allotropes

Iron represents perhaps the best-known example of allotropy in a metal. There are three allotropic forms of iron, known as alpha, gamma, and delta.

As molten iron cools down it crystallizes at 1538 °C into its delta allotrope, which has a body-centered cubic (BCC) crystal structure. As it cools further its crystal structure changes to face-centered cubic (FCC) at 1394 °C, when it is known as gamma-iron, or austenite. At 912 °C the crystal structure again becomes BCC as alpha-iron, or ferrite, is formed, and at 770 °C (the Curie point, Tc) the iron becomes magnetic. As the iron passes through the Curie temperature there is no change in crystalline structure, but there is a change in "domain structure", where each domain contains iron atoms with a particular electronic spin. In unmagnetized iron, all the electronic spins of the atoms within one domain are in the same direction; however, in neighbouring domains they point in various directions and thus cancel out. In magnetized iron, the electronic spins of all the domains are all aligned, so that the magnetic effects of neighbouring domains reinforce each other. Although each domain contains billions of atoms, they are very small, about one thousandth of a centimetre across.

Iron is of most importance when mixed with certain other metals and with carbon to form steels. There are many types of steels, all with different properties; and an understanding of the properties of the allotropes of iron is key to the manufacture of good quality steels.

Alpha iron, also known as ferrite, is the most stable form of iron at normal temperatures. It is a fairly soft metal that can dissolve only a small concentration of carbon (no more than 0.021% by mass at 910 °C).

Above 912 °C and up to 1401 °C alpha iron undergoes a phase transition from body-centred cubic to the face-centred cubic configuration of gamma iron, also called austenite. This is similarly soft and metallic but can dissolve considerably more carbon (as much as 2.04% by mass at 1146 °C). This form of iron is used in the type of stainless steel used for making cutlery, and hospital and food-service equipment.

Applications

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Iron is the most widely used of all the metals, accounting for 95% of worldwide metal production. Its low cost and high strength make it indispensable in engineering applications such as the construction of machinery and machine tools, automobiles, the hulls of large ships, and structural components for buildings. Since pure iron is quite soft, it is most commonly used in the form of steel. Some of the forms in which iron is produced commercially include:

• Pig iron has 3.5–4.5% carbon^[5] and contains varying amounts of contaminants such as sulfur, silicon and phosphorus. Its only significance is that of an intermediate step on the way from iron ore to cast iron and steel.
• Cast iron contains 2–4% carbon, 1–6% silicon, and small amounts of manganese. Contaminants present in pig iron that negatively affect material properties, such as sulfur and phosphorus, have been reduced to an acceptable level. It has a melting point in the range of 1420–1470 K, which is lower than either of its two main components, and makes it the first product to be melted when carbon and iron are heated together. Its mechanical properties vary greatly, dependent upon the form carbon takes in the alloy. "White" cast irons contain their carbon in the form of cementite, or iron carbide. This hard, brittle compound dominates the mechanical properties of white cast irons, rendering them hard, but unresistant to shock. The broken surface of a white cast iron is full of fine facets of the broken carbide, a very pale, silvery, shiny material, hence the appellation. In grey iron the carbon exists free as fine flakes of graphite, and also renders the material brittle due to the stress-raising nature of the sharp edged flakes of graphite. A newer variant of grey iron, referred to as ductile iron is specially treated with trace amounts of magnesium to alter the shape of graphite to spheroids, or nodules, vastly increasing the toughness and strength of the material.
  

  

  The fining process of smelting iron ore to make wrought iron from pig iron, with the right illustration displaying men working a blast furnace, from the Tiangong Kaiwu encyclopedia, published in 1637 by Song Yingxing.

• Wrought iron contains less than 0.25% carbon.^[5] It is a tough, malleable product, but not as fusible as pig iron. If honed to an edge, it loses it quickly. Wrought iron is characterized by the presence of fine fibers of slag entrapped in the metal. Wrought iron is more corrosion resistant than steel. It has been almost completely replaced by mild steel for traditional "wrought iron" products and blacksmithing. Mild steel corrodes more readily that wrought iron, but is cheaper and more widely available.
• Carbon steel contains 2.0% carbon or less,^[6] with small amounts of manganese, sulfur, phosphorus, and silicon.
• Alloy steels contain varying amounts of carbon as well as other metals, such as chromium, vanadium, molybdenum, nickel, tungsten, etc. Their alloy content raises their cost, and so they can usually only be justified for specialist uses. Recent developments in ferrous metallurgy have produced a growing range of microalloyed steels, also termed 'HSLA' or high-strength, low alloy steels, containing tiny additions to produce high strengths and often spectacular toughness at minimal cost.
• Iron(III) oxides are used in the production of magnetic storage media in computers. They are often mixed with other compounds, and retain their magnetic properties in solution.

The main disadvantage of iron and steel is that pure iron, and most of its alloys, suffer badly from rust if not protected in some way. Painting, galvanization, passivation, plastic coating and bluing are some techniques used to protect iron from rust by excluding water and oxygen or by sacrificial protection.

Iron compounds

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• Iron oxides (FeO, Fe₃O₄, and Fe₂O₃) are ores used for iron production (see bloomery and blast furnace). They are also used as a catalyst in the Space Shuttle Solid Rocket Boosters.^[7]
• Iron(II) acetate (Fe(CH₃CO₂)₂ (ferrous acetate) is used as a mordant in the dyeing of cloth and leather, and as a wood preservative.
• Iron(III) ammonium citrate (C₆H_5+4yFe_xN_yO₇) is used in blueprints.
• Iron(III) arsenate (FeAsO₄) is used in insecticides.
• Iron(III) chloride (FeCl₃) is used: in water purification and sewage treatment, in the dyeing of cloth, as a coloring agent in paints, as an additive in animal feed, and as an etchant for copper in the manufacture of printed circuits.
• Iron(III) chromate (Fe₂(CrO₄)₃) is used as a yellow pigment for paints and ceramics.
• Iron(III) hydroxide (Fe(OH)₃) is used as a brown pigment for rubber and in water purification systems.
• Iron(III) phosphate (FePO₄) is used in fertilizer and as an additive in human and animal food.
• Iron(II) gluconate (Fe(C₆H₁₁O₇)₂) is used as a dietary supplement in iron pills.
• Iron(II) oxalate (FeC₂O₄) is used as yellow pigment for paints, plastics, glass and ceramics, and in photography.
• Iron(II) sulfate (FeSO₄) is used in water purification and sewage treatment systems, as a catalyst in the production of ammonia, as an ingredient in fertilizer, herbicide, and moss killer, as an additive in animal feed, in wood preservative, and as an additive to flour to increase nutritional iron levels. Experimental iron fertilization of areas of the ocean using iron(II) sulfate has proven successful in increasing plankton growth.^[8][9][10]
• Iron-fluorine complex (FeF₆)^3- is found in solutions containing both Fe(III) ions and fluoride ions.
• Potassium ferrate (K₂FeO₄) contains a tetra-oxo anion of iron in its highest known oxidation State.

History

This section requires expansion.

The first wrought iron used by mankind during prehistory came from meteors. The smelting of iron in bloomeries began in the second millennium BC. Artifacts from smelted iron occur in India from 1800-1200 BC.^[11] in the Levant from about 1500 BC (suggesting smelting in Anatolia or the Caucasus).^[12][13]

Cast iron was first produced in China about 550 BC,^[14] but not in Europe until the medieval period. During the medieval period, means were found in Europe of producing wrought iron from cast iron (in this context known as pig iron) using finery forges. For all these processes, charcoal was required as fuel.

Steel (with smaller carbon content than pig iron but more than wrought iron) was first produced in antiquity. New methods of producing it by carburizing bars of iron in the cementation process were devised in the 17th century AD. In the Industrial Revolution, new methods of producing bar iron without charcoal were devised and these were later applied to produce steel. In the late 1850s, Henry Bessemer invented a new steelmaking process, involving blowing air through molten pig iron, to produce mild steel. This and other 19th century and later processes have led to wrought iron no longer being produced.

Production of iron from iron ore

The production of iron or steel is a process unless the desired final product is cast iron. The first stage is to produce pig iron in a blast furnace. The second is to make wrought iron or steel from pig iron by a further process.

Blast furnace

Ninety percent of all mining of metallic ores is for the extraction of iron. Industrially, iron is produced starting from iron ores, principally hematite (nominally Fe₂O₃) and magnetite (Fe₃O₄) by a carbothermic reaction (reduction with carbon) in a blast furnace at temperatures of about 2000 °C. In a blast furnace, iron ore, carbon in the form of coke, and a flux such as limestone (which is used to remove impurities in the ore which would otherwise clog the furnace with solid material) are fed into the top of the furnace, while a blast of heated air is forced into the furnace at the bottom.

In the furnace, the coke reacts with oxygen in the air blast to produce carbon monoxide:

2 C + O₂ → 2 CO

The carbon monoxide reduces the iron ore (in the chemical equation below, hematite) to molten iron, becoming carbon dioxide in the process:

3 CO + Fe₂O₃ → 2 Fe + 3 CO₂

The flux is present to melt impurities in the ore, principally silicon dioxide sand and other silicates. Common fluxes include limestone (principally calcium carbonate) and dolomite (calcium-magnesium carbonate). Other fluxes may be used depending on the impurities that need to be removed from the ore. In the heat of the furnace the limestone flux decomposes to calcium oxide (quicklime):

CaCO₃ → CaO + CO₂

Then calcium oxide combines with silicon dioxide to form a slag.

CaO + SiO₂ → CaSiO₃

The slag melts in the heat of the furnace. In the bottom of the furnace, the molten slag floats on top of the denser molten iron, and apertures in the side of the furnace are opened to run off the iron and the slag separately. The iron once cooled, is called pig iron, while the slag can be used as a material in road construction or to improve mineral-poor soils for agriculture.^{[citation needed]}

In 2005, approximately 1,544 Mt (million metric tons) of iron ore were produced worldwide. China was the top producer of iron ore with at least one-fourth world share followed by Brazil, Australia and India, reports the British Geological Survey.

Further processes

Pig iron is not pure iron, but has 4-5% carbon dissolved in it with small amounts of other impurities like sulfur, magnesium, phosphorus and manganese. As the carbon is the major impurity, the iron (pig iron) becomes brittle and hard. This form of iron is used to cast articles in foundries such as stoves, pipes, radiators, lamp-posts and rails.

Alternatively pig iron may be made into steel (with up to about 2% carbon) or wrought iron (commercially pure iron). Various processes have been used for this, including finery forges, puddling furnaces, Bessemer converters, open hearth furnaces, basic oxygen furnaces, and electric arc furnaces. In all cases, the objective is to oxidize some or all of the carbon, together with other impurities. On the other hand, other metals may be added to make alloy steels.

The hardness of the steel depends upon its carbon content, the higher the proportion of carbon, the greater the hardness and the lesser the ductility. The properties of the steel can also be changed by tempering it. To harden the steel, it is heated to red hot and then cooled by quenching it in the water. It becomes harder and more brittle. This steel is then heated to a required temperature and allowed to cool. The steel thus formed is less brittle.

Isotopes

Naturally occurring iron consists of four isotopes: 5.845% of radioactive ⁵⁴Fe (half-life: >3.1×10²² years), 91.754% of stable ⁵⁶Fe, 2.119% of stable ⁵⁷Fe and 0.282% of stable ⁵⁸Fe. ⁶⁰Fe is an extinct radionuclide of long half-life (1.5 million years).

Much of the past work on measuring the isotopic composition of Fe has centered on determining ⁶⁰Fe variations due to processes accompanying nucleosynthesis (i.e., meteorite studies) and ore formation. In the last decade however, advances in mass spectrometry technology have allowed the detection and quantification of minute, naturally occurring variations in the ratios of the stable isotopes of iron. Much of this work has been driven by the Earth and planetary science communities, although applications to biological and industrial systems are beginning to emerge.^[15]

The most abundant iron isotope ⁵⁶Fe is of particular interest to nuclear scientists. A common misconception is that this isotope represents the most stable nucleus possible, and that it thus would be impossible to perform fission or fusion on ⁵⁶Fe and still liberate energy. This is not true, as both ⁶²Ni and ⁵⁸Fe are more stable, being the most stable nuclei. However, since ⁵⁶Ni is much more easily produced from lighter nuclei in the alpha process in nuclear reactions in supernovae (see silicon burning process), nickel-56 (14 alpha particles) is the endpoint of fusion chains inside extremely massive stars, since addition of another alpha would result in zinc-60, which requires a great deal more energy. This nickel-56, which has a half-life of about 6 days, is therefore made in quantity in these stars, but soon decays by two successive positron emissions within supernova decay products in the supernova remnant gas cloud, to first radioactive cobalt-56, and then stable iron-56. This last nuclide is therefore common in the universe, relative to other stable metals of approximately the same atomic weight.

In phases of the meteorites Semarkona and Chervony Kut a correlation between the concentration of ⁶⁰Ni, the daughter product of ⁶⁰Fe, and the abundance of the stable iron isotopes could be found which is evidence for the existence of ⁶⁰Fe at the time of formation of the solar system. Possibly the energy released by the decay of ⁶⁰Fe contributed, together with the energy released by decay of the radionuclide ²⁶Al, to the remelting and differentiation of asteroids after their formation 4.6 billion years ago. The abundance of ⁶⁰Ni present in extraterrestrial material may also provide further insight into the origin of the solar system and its early history. Of the stable isotopes, only ⁵⁷Fe has a nuclear spin (−1/2).

Iron in organic synthesis

The use of iron metal filings in organic synthesis is mainly for the reduction of nitro compounds.^[16] Additionally, iron has been used for desulfurizations,^[17] reduction of aldehydes,^[18] and the deoxygenation of amine oxides.^[19]

Iron in biology

Iron is essential to nearly all known organisms. In cells, iron is generally stored in the centre of metalloproteins, because "free" iron (which binds non-specifically to many cellular components) can catalyse production of toxic free radicals. Iron deficiency can lead to iron deficiency anemia.

In animals, plants, and fungi, iron is often the metal ion incorporated into the heme complex. Heme is an essential component of cytochrome proteins, which mediate redox reactions, and of oxygen carrier proteins such as hemoglobin, myoglobin, and leghemoglobin. Inorganic iron also contributes to redox reactions in the iron-sulfur clusters of many enzymes, such as nitrogenase (involved in the synthesis of ammonia from nitrogen and hydrogen) and hydrogenase. Non-heme iron proteins include the enzymes methane monooxygenase (oxidizes methane to methanol), ribonucleotide reductase (reduces ribose to deoxyribose; DNA biosynthesis), hemerythrins (oxygen transport and fixation in marine invertebrates) and purple acid phosphatase (hydrolysis of phosphate esters).

Iron distribution is heavily regulated in mammals, partly because iron has a high potential for biological toxicity. Iron distribution is also regulated because many bacteria require iron, so restricting its availability to bacteria (generally by sequestering it inside cells) can help to prevent or limit infections. This is probably the reason for the relatively low amounts of iron in mammalian milk. A major component of this regulation is the protein transferrin, which binds iron absorbed from the duodenum and carries it in the blood to cells.^[20]

Nutrition and dietary sources

Good sources of dietary iron include red meat, fish, poultry, lentils, beans, leaf vegetables, tofu, chickpeas, black-eyed peas, fortified bread, and fortified breakfast cereals. Iron in low amounts is found in molasses, teff and farina. Iron in meat (haem iron) is more easily absorbed than iron in vegetables,^[21] but heme/hemoglobin from red meat has effects which may increase the likelihood of colorectal cancer.^[22][23]

Iron provided by dietary supplements is often found as iron (II) fumarate, although iron sulfate is cheaper and is absorbed equally well. Elemental iron, despite being absorbed to a much smaller extent (stomach acid is sufficient to convert some of it to ferrous iron), is often added to foods such as breakfast cereals or "enriched" wheat flour (where it is listed as "reduced iron" in the list of ingredients). Iron is most available to the body when chelated to amino acids - iron in this form is ten to fifteen times more bioavailable^[24] than any other, and is also available for use as a common iron supplement. Often the amino acid chosen for this purpose is the cheapest and most common amino acid, glycine, leading to "iron glycinate" supplements.^[25] The RDA for iron varies considerably based on age, gender, and source of dietary iron (heme-based iron has higher bioavailability).^[26] Infants may require iron supplements if they are breast-fed.^[27] Blood donors and pregnant women are at special risk of low iron levels and are often advised to supplement their iron intake.^{[citation needed]}

Regulation of iron uptake

Excessive iron can be toxic, because free ferrous iron reacts with peroxides to produce free radicals, which are highly reactive and can damage DNA, proteins, lipids, and other cellular components. Thus, iron toxicity occurs when there is free iron in the cell, which generally occurs when iron levels exceed the capacity of transferrin to bind the iron.

Iron uptake is tightly regulated by the human body, which has no regulated physiological means of excreting iron. Only small amounts of iron are lost daily due to mucosal and skin epithelial cell sloughing, so control of iron levels is mostly by regulating uptake.^[28] However, large amounts of ingested iron can cause excessive levels of iron in the blood because high iron levels can damage the cells of the gastrointestinal tract, preventing them from regulating iron absorption, high blood concentrations of iron damage cells in the heart, liver and elsewhere, which can cause serious problems, including long-term organ damage and even death.

Humans experience iron toxicity above 20 milligrams of iron for every kilogram of mass, and 60 milligrams per kilogram is a lethal dose.^[29] Over-consumption of iron, often the result of children eating large quantities of ferrous sulfate tablets intended for adult consumption, is one of the most common toxicological causes of death in children under six.^[29] The DRI lists the Tolerable Upper Intake Level (UL) for adults as 45 mg/day. For children under fourteen years old the UL is 40 mg/day.

Regulation of iron uptake is impaired in some people as a result of a genetic defect that maps to the HLA-H gene region on chromosome 6. In these people, excessive iron intake can result in iron overload disorders, such as hemochromatosis. Many people have a genetic susceptibility to iron overload without realizing it or being aware of a family history of the problem. For this reason, it is advised that people do not take iron supplements unless they suffer from iron deficiency and have consulted a doctor. Hemochromatosis is estimated to cause disease in between 0.3 and 0.8% of Caucasians.^[30]

The medical management of iron toxicity is complex, and can include use of a specific chelating agent called deferoxamine to bind and expel excess iron from the body.

See also

• Al-Hadid (Iron) in the Qur'an.
• El Mutún in Bolivia, where 20% of the world's accessible iron and magnesium is located.
• Iron Age
• Iron fertilization - Proposed fertilization of oceans to stimulate phytoplankton growth.
• Iron (metaphor)
• Iron in mythology
• List of countries by iron production
• Pelletizing - Process of creation of iron ore pellets.
• Specht Building - A historic landmark in Omaha, Nebraska utilizing an iron facade.

Bibliography

• Doulias PT, Christoforidis S, Brunk UT, Galaris D. Endosomal and lysosomal effects of desferrioxamine: protection of HeLa cells from hydrogen peroxide-induced DNA damage and induction of cell-cycle arrest. Free Radic Biol Med. 2003;35:719-28.
• H. R. Schubert, History of the British Iron and Steel Industry ... to 1775 AD (Routledge, London, 1957)
• R. F. Tylecote, History of Metallurgy (Institute of Materials, London 1992).
• R. F. Tylecote, 'Iron in the Industrial Revolution' in J. Day and R. F. Tylecote, The Industrial Revolution in Metals (Institute of Materials 1991), 200-60.
• Los Alamos National Laboratory — Iron
• Crystal structure of iron

References

External links

Wikimedia Commons has media related to: Iron

• WebElements.com – Iron
• It's Elemental – Iron
• The Most Tightly Bound Nuclei

v • d • e Periodic table
H																															He
Li	Be																									B	C	N	O	F	Ne
Na	Mg																									Al	Si	P	S	Cl	Ar
K	Ca															Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn	Ga	Ge	As	Se	Br	Kr
Rb	Sr															Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd	In	Sn	Sb	Te	I	Xe
Cs	Ba	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg	Tl	Pb	Bi	Po	At	Rn
Fr	Ra	Ac	Th	Pa	U	Np	Pu	Am	Cm	Bk	Cf	Es	Fm	Md	No	Lr	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Uub	Uut	Uuq	Uup	Uuh	Uus	Uuo
Alkali metals Alkaline earth metals Lanthanoids Actinoids Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases

v • d • e   Iron compounds FeBr2 · FeBr3 · FeCl2 · FeCl3 · FeF2 · FeF3 · Fe(NO3)3 · FeO · FeO(OH) · Fe(OH)2 · FePO4 · FeS · FeSO4 · Fe2(CrO4)3 · Fe2O3 · Fe2(SO4)3 · Fe3O4 · Fe3(PO4)2