Iron

This article is about the metallic element. For other uses, see Iron (disambiguation).

Iron is a chemical element; it has the symbol Fe (from Latin ferrum 'iron') and atomic number 26. It is a metal that belongs to the first transition series and group 8 of the periodic table. It is, by mass, the most common element on Earth, forming much of Earth's outer and inner core. It is the fourth most abundant element in the Earth's crust, being mainly deposited by meteorites in its metallic state.

Iron, ₂₆Fe

Iron
Pronunciation	/ˈaɪərn/
Allotropes	see Allotropes of iron
Appearance	lustrous metallic with a grayish tinge
Standard atomic weight Ar°(Fe)
	55.845±0.002[1] 55.845±0.002 (abridged)[2]
Iron in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson – ↑ Fe ↓ Ru manganese ← iron → cobalt
Atomic number (Z)	26
Group	group 8
Period	period 4
Block	d-block
Electron configuration	[Ar] 3d6 4s2
Electrons per shell	2, 8, 14, 2
Physical properties
Phase at STP	solid
Melting point	1811 K ​(1538 °C, ​2800 °F)
Boiling point	3134 K ​(2861 °C, ​5182 °F)
Density (at 20° C)	7.874 g/cm3 [3]
when liquid (at m.p.)	6.98 g/cm3
Heat of fusion	13.81 kJ/mol
Heat of vaporization	340 kJ/mol
Molar heat capacity	25.10 J/(mol·K)
Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 1728 1890 2091 2346 2679 3132
Atomic properties
Oxidation states	−4, −2, −1, 0, +1,[4] +2, +3, +4, +5,[5] +6, +7[6] (an amphoteric oxide)
Electronegativity	Pauling scale: 1.83
Ionization energies	1st: 762.5 kJ/mol 2nd: 1561.9 kJ/mol 3rd: 2957 kJ/mol (more)
Atomic radius	empirical: 126 pm
Covalent radius	Low spin: 132±3 pm High spin: 152±6 pm
Van der Waals radius	194  pm
Spectral lines of iron
Other properties
Natural occurrence	primordial
Crystal structure	α-Fe: ​body-centered cubic (bcc) (cI2)
Lattice constant	a = 286.65 pm (at 20 °C)[3]
Crystal structure	γ-Fe (912–1394 °C): ​face-centered cubic (fcc) (cF4)
Lattice constant	a = 364.68 pm (at 916 °C)[7]
Thermal expansion	12.07×10−6/K (at 20 °C)[3]
Thermal conductivity	80.4 W/(m⋅K)
Electrical resistivity	96.1 nΩ⋅m (at 20 °C)
Curie point	1043 K
Magnetic ordering	ferromagnetic
Young's modulus	211 GPa
Shear modulus	82 GPa
Bulk modulus	170 GPa
Speed of sound thin rod	5120 m/s (at r.t.) (electrolytic)
Poisson ratio	0.29
Mohs hardness	4
Vickers hardness	608 MPa
Brinell hardness	200–1180 MPa
CAS Number	7439-89-6
History
Discovery	before 5000 BC
Symbol	"Fe": from Latin ferrum
Isotopes of iron v e
Main isotopes[8] Decay abun­dance half-life (t1/2) mode pro­duct 54Fe 5.85% stable 55Fe synth 2.73 y ε 55Mn 56Fe 91.8% stable 57Fe 2.12% stable 58Fe 0.28% stable 59Fe synth 44.6 d β− 59Co 60Fe trace 2.6×106 y β− 60Co
Category: Iron view talk edit | references

Extracting usable metal from iron ores requires kilns or furnaces capable of reaching 1,500 °C (2,730 °F), about 500 °C (932 °F) higher than that required to smelt copper. Humans started to master that process in Eurasia during the 2nd millennium BC and the use of iron tools and weapons began to displace copper alloys – in some regions, only around 1200 BC. That event is considered the transition from the Bronze Age to the Iron Age. In the modern world, iron alloys, such as steel, stainless steel, cast iron and special steels, are by far the most common industrial metals, due to their mechanical properties and low cost. The iron and steel industry is thus very important economically, and iron is the cheapest metal, with a price of a few dollars per kilogram or pound.

Pristine and smooth pure iron surfaces are a mirror-like silvery-gray. Iron reacts readily with oxygen and water to produce brown-to-black hydrated iron oxides, commonly known as rust. Unlike the oxides of some other metals that form passivating layers, rust occupies more volume than the metal and thus flakes off, exposing more fresh surfaces for corrosion. Chemically, the most common oxidation states of iron are iron(II) and iron(III). Iron shares many properties of other transition metals, including the other group 8 elements, ruthenium and osmium. Iron forms compounds in a wide range of oxidation states, −4 to +7. Iron also forms many coordination compounds; some of them, such as ferrocene, ferrioxalate, and Prussian blue have substantial industrial, medical, or research applications.

The body of an adult human contains about 4 grams (0.005% body weight) of iron, mostly in hemoglobin and myoglobin. These two proteins play essential roles in oxygen transport by blood and oxygen storage in muscles. To maintain the necessary levels, human iron metabolism requires a minimum of iron in the diet. Iron is also the metal at the active site of many important redox enzymes dealing with cellular respiration and oxidation and reduction in plants and animals.^[9]

Allotropes

Main article: Allotropes of iron

Molar volume vs. pressure for α iron at room temperature

At least four allotropes of iron (differing atom arrangements in the solid) are known, conventionally denoted α, γ, δ, and ε.

The first three forms are observed at ordinary pressures. As molten iron cools past its freezing point of 1538 °C, it crystallizes into its δ allotrope, which has a body-centered cubic (bcc) crystal structure. As it cools further to 1394 °C, it changes to its γ-iron allotrope, a face-centered cubic (fcc) crystal structure, or austenite. At 912 °C and below, the crystal structure again becomes the bcc α-iron allotrope.^[10]

The physical properties of iron at very high pressures and temperatures have also been studied extensively,^[11][12] because of their relevance to theories about the cores of the Earth and other planets. Above approximately 10 GPa and temperatures of a few hundred kelvin or less, α-iron changes into another hexagonal close-packed (hcp) structure, which is also known as ε-iron. The higher-temperature γ-phase also changes into ε-iron, but does so at higher pressure.

Some controversial experimental evidence exists for a stable β phase at pressures above 50 GPa and temperatures of at least 1500 K. It is supposed to have an orthorhombic or a double hcp structure.^[13] (Confusingly, the term "β-iron" is sometimes also used to refer to α-iron above its Curie point, when it changes from being ferromagnetic to paramagnetic, even though its crystal structure has not changed.^[10])

The inner core of the Earth is generally presumed to consist of an iron-nickel alloy with ε (or β) structure.^[14]

Melting and boiling points

Low-pressure phase diagram of pure iron

The melting and boiling points of iron, along with its enthalpy of atomization, are lower than those of the earlier 3d elements from scandium to chromium, showing the lessened contribution of the 3d electrons to metallic bonding as they are attracted more and more into the inert core by the nucleus;^[15] however, they are higher than the values for the previous element manganese because that element has a half-filled 3d sub-shell and consequently its d-electrons are not easily delocalized. This same trend appears for ruthenium but not osmium.^[16]

The melting point of iron is experimentally well defined for pressures less than 50 GPa. For greater pressures, published data (as of 2007) still varies by tens of gigapascals and over a thousand kelvin.^[17]

Magnetic properties

Magnetization curves of 9 ferromagnetic materials, showing saturation. 1. Sheet steel, 2. Silicon steel, 3. Cast steel, 4. Tungsten steel, 5. Magnet steel, 6. Cast iron, 7. Nickel, 8. Cobalt, 9. Magnetite^[18]

Below its Curie point of 770 °C (1,420 °F; 1,040 K), α-iron changes from paramagnetic to ferromagnetic: the spins of the two unpaired electrons in each atom generally align with the spins of its neighbors, creating an overall magnetic field.^[19] This happens because the orbitals of those two electrons (d_z² and d_{x² − y²}) do not point toward neighboring atoms in the lattice, and therefore are not involved in metallic bonding.^[10]

In the absence of an external source of magnetic field, the atoms get spontaneously partitioned into magnetic domains, about 10 micrometers across,^[20] such that the atoms in each domain have parallel spins, but some domains have other orientations. Thus a macroscopic piece of iron will have a nearly zero overall magnetic field.

Application of an external magnetic field causes the domains that are magnetized in the same general direction to grow at the expense of adjacent ones that point in other directions, reinforcing the external field. This effect is exploited in devices that need to channel magnetic fields to fulfill design function, such as electrical transformers, magnetic recording heads, and electric motors. Impurities, lattice defects, or grain and particle boundaries can "pin" the domains in the new positions, so that the effect persists even after the external field is removed – thus turning the iron object into a (permanent) magnet.^[19]

Similar behavior is exhibited by some iron compounds, such as the ferrites including the mineral magnetite, a crystalline form of the mixed iron(II,III) oxide Fe₃O₄ (although the atomic-scale mechanism, ferrimagnetism, is somewhat different). Pieces of magnetite with natural permanent magnetization (lodestones) provided the earliest compasses for navigation. Particles of magnetite were extensively used in magnetic recording media such as core memories, magnetic tapes, floppies, and disks, until they were replaced by cobalt-based materials.

Isotopes

Main article: Isotopes of iron

Iron has four stable isotopes: ⁵⁴Fe (5.845% of natural iron), ⁵⁶Fe (91.754%), ⁵⁷Fe (2.119%) and ⁵⁸Fe (0.282%). Twenty-four artificial isotopes have also been created. Of these stable isotopes, only ⁵⁷Fe has a nuclear spin (−1⁄2). The nuclide ⁵⁴Fe theoretically can undergo double electron capture to ⁵⁴Cr, but the process has never been observed and only a lower limit on the half-life of 4.4×10²⁰ years has been established.^[21]

⁶⁰Fe is an extinct radionuclide of long half-life (2.6 million years).^[22] It is not found on Earth, but its ultimate decay product is its granddaughter, the stable nuclide ⁶⁰Ni.^[8] Much of the past work on isotopic composition of iron has focused on the nucleosynthesis of ⁶⁰Fe through studies of meteorites and ore formation. In the last decade, advances in mass spectrometry have allowed the detection and quantification of minute, naturally occurring variations in the ratios of the stable isotopes of iron. Much of this work is driven by the Earth and planetary science communities, although applications to biological and industrial systems are emerging.^[23]

In phases of the meteorites Semarkona and Chervony Kut, a correlation between the concentration of ⁶⁰Ni, the granddaughter of ⁶⁰Fe, and the abundance of the stable iron isotopes provided evidence for the existence of ⁶⁰Fe at the time of formation of the Solar System. Possibly the energy released by the decay of ⁶⁰Fe, along with that released by ²⁶Al, contributed to the remelting and differentiation of asteroids after their formation 4.6 billion years ago. The abundance of ⁶⁰Ni present in extraterrestrial material may bring further insight into the origin and early history of the Solar System.^[24]

The most abundant iron isotope ⁵⁶Fe is of particular interest to nuclear scientists because it represents the most common endpoint of nucleosynthesis.^[25] Since ⁵⁶Ni (14 alpha particles) is easily produced from lighter nuclei in the alpha process in nuclear reactions in supernovae (see silicon burning process), it is the endpoint of fusion chains inside extremely massive stars. Although adding more alpha particles is possible, but nonetheless the sequence does effectively end at ⁵⁶Ni because conditions in stellar interiors cause the competition between photodisintegration and the alpha process to favor photodisintegration around ⁵⁶Ni.^[26][27] This ⁵⁶Ni, which has a half-life of about 6 days, is created in quantity in these stars, but soon decays by two successive positron emissions within supernova decay products in the supernova remnant gas cloud, first to radioactive ⁵⁶Co, and then to stable ⁵⁶Fe. As such, 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.^[28] It is also very common in the universe, relative to other stable metals of approximately the same atomic weight.^[28][29] Iron is the sixth most abundant element in the universe, and the most common refractory element.^[30]

Photon mass attenuation coefficient for iron

Although a further tiny energy gain could be extracted by synthesizing ⁶²Ni, which has a marginally higher binding energy than ⁵⁶Fe, conditions in stars are unsuitable for this process. Element production in supernovas greatly favor iron over nickel, and in any case, ⁵⁶Fe still has a lower mass per nucleon than ⁶²Ni due to its higher fraction of lighter protons.^[31] Hence, elements heavier than iron require a supernova for their formation, involving rapid neutron capture by starting ⁵⁶Fe nuclei.^[28]

In the far future of the universe, assuming that proton decay does not occur, cold fusion occurring via quantum tunnelling would cause the light nuclei in ordinary matter to fuse into ⁵⁶Fe nuclei. Fission and alpha-particle emission would then make heavy nuclei decay into iron, converting all stellar-mass objects to cold spheres of pure iron.^[32]

Cosmogenesis

Iron's abundance in rocky planets like Earth is due to its abundant production during the runaway fusion and explosion of type Ia supernovae, which scatters the iron into space.^[33][34]

Metallic iron

A polished and chemically etched piece of an iron meteorite, believed to be similar in composition to the Earth's metallic core, showing individual crystals of the iron-nickel alloy (Widmanstatten pattern)

Metallic or native iron is rarely found on the surface of the Earth because it tends to oxidize. However, both the Earth's inner and outer core, which together account for 35% of the mass of the whole Earth, are believed to consist largely of an iron alloy, possibly with nickel. Electric currents in the liquid outer core are believed to be the origin of the Earth's magnetic field. The other terrestrial planets (Mercury, Venus, and Mars) as well as the Moon are believed to have a metallic core consisting mostly of iron. The M-type asteroids are also believed to be partly or mostly made of metallic iron alloy.

The rare iron meteorites are the main form of natural metallic iron on the Earth's surface. Items made of cold-worked meteoritic iron have been found in various archaeological sites dating from a time when iron smelting had not yet been developed; and the Inuit in Greenland have been reported to use iron from the Cape York meteorite for tools and hunting weapons.^[35] About 1 in 20 meteorites consist of the unique iron-nickel minerals taenite (35–80% iron) and kamacite (90–95% iron).^[36] Native iron is also rarely found in basalts that have formed from magmas that have come into contact with carbon-rich sedimentary rocks, which have reduced the oxygen fugacity sufficiently for iron to crystallize. This is known as telluric iron and is described from a few localities, such as Disko Island in West Greenland, Yakutia in Russia and Bühl in Germany.^[37]

Mantle minerals

Ferropericlase (Mg,Fe)O, a solid solution of periclase (MgO) and wüstite (FeO), makes up about 20% of the volume of the lower mantle of the Earth, which makes it the second most abundant mineral phase in that region after silicate perovskite (Mg,Fe)SiO₃; it also is the major host for iron in the lower mantle.^[38] At the bottom of the transition zone of the mantle, the reaction γ-(Mg,Fe)₂[SiO₄] ↔ (Mg,Fe)[SiO₃] + (Mg,Fe)O transforms γ-olivine into a mixture of silicate perovskite and ferropericlase and vice versa. In the literature, this mineral phase of the lower mantle is also often called magnesiowüstite.^[39] Silicate perovskite may form up to 93% of the lower mantle,^[40] and the magnesium iron form, (Mg,Fe)SiO₃, is considered to be the most abundant mineral in the Earth, making up 38% of its volume.^[41]

Earth's crust

Ochre path in Roussillon

While iron is the most abundant element on Earth, most of this iron is concentrated in the inner and outer cores.^[42][43] The fraction of iron that is in Earth's crust only amounts to about 5% of the overall mass of the crust and is thus only the fourth most abundant element in that layer (after oxygen, silicon, and aluminium).^[44]

Most of the iron in the crust is combined with various other elements to form many iron minerals. An important class is the iron oxide minerals such as hematite (Fe₂O₃), magnetite (Fe₃O₄), and siderite (FeCO₃), which are the major ores of iron. Many igneous rocks also contain the sulfide minerals pyrrhotite and pentlandite.^[45][46] During weathering, iron tends to leach from sulfide deposits as the sulfate and from silicate deposits as the bicarbonate. Both of these are oxidized in aqueous solution and precipitate in even mildly elevated pH as iron(III) oxide.^[47]

Banded iron formation in McKinley Park, Minnesota

Large deposits of iron are banded iron formations, a type of rock consisting of repeated thin layers of iron oxides alternating with bands of iron-poor shale and chert. The banded iron formations were laid down in the time between 3,700 million years ago and 1,800 million years ago.^[48][49]

Materials containing finely ground iron(III) oxides or oxide-hydroxides, such as ochre, have been used as yellow, red, and brown pigments since pre-historical times. They contribute as well to the color of various rocks and clays, including entire geological formations like the Painted Hills in Oregon and the Buntsandstein ("colored sandstone", British Bunter).^[50] Through Eisensandstein (a jurassic 'iron sandstone', e.g. from Donzdorf in Germany)^[51] and Bath stone in the UK, iron compounds are responsible for the yellowish color of many historical buildings and sculptures.^[52] The proverbial red color of the surface of Mars is derived from an iron oxide-rich regolith.^[53]

Significant amounts of iron occur in the iron sulfide mineral pyrite (FeS₂), but it is difficult to extract iron from it and it is therefore not exploited.^[54] In fact, iron is so common that production generally focuses only on ores with very high quantities of it.^[55]

According to the International Resource Panel's Metal Stocks in Society report, the global stock of iron in use in society is 2,200 kg per capita. More-developed countries differ in this respect from less-developed countries (7,000–14,000 vs 2,000 kg per capita).^[56]

Oceans

Ocean science demonstrated the role of the iron in the ancient seas in both marine biota and climate.^[57]

Main article: Iron compounds

Oxidation state	Representative compound
−4 (d10s2)	[FeIn6−xSnx][58]
−2 (d10)	Disodium tetracarbonylferrate (Collman's reagent)
−1 (d9)	Fe 2(CO)2− 8
0 (d8)	Iron pentacarbonyl
1 (d7)	Cyclopentadienyliron dicarbonyl dimer ("Fp2")
2 (d6)	Ferrous sulfate, Ferrocene
3 (d5)	Ferric chloride, Ferrocenium tetrafluoroborate
4 (d4)	Fe(diars) 2Cl2+ 2, FeO(BF4)2
5 (d3)	FeO3− 4
6 (d2)	Potassium ferrate
7 (d1)	[FeO4]– (matrix isolation, 4K)

Iron shows the characteristic chemical properties of the transition metals, namely the ability to form variable oxidation states differing by steps of one and a very large coordination and organometallic chemistry: indeed, it was the discovery of an iron compound, ferrocene, that revolutionalized the latter field in the 1950s.^[59] Iron is sometimes considered as a prototype for the entire block of transition metals, due to its abundance and the immense role it has played in the technological progress of humanity.^[60] Its 26 electrons are arranged in the configuration [Ar]3d⁶4s², of which the 3d and 4s electrons are relatively close in energy, and thus a number of electrons can be ionized.