Allotropes

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.^[7]

The physical properties of iron at very high pressures and temperatures have also been studied extensively,^[8][9] 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.^[10] (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.^[7])

The inner core of the Earth is generally presumed to consist of an iron-nickel alloy with ε (or β) structure.^[11]

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;^[12] 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.^[13]

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.^[14]

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^[15]

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.^[16] 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.^[7]

In the absence of an external source of magnetic field, the atoms get spontaneously partitioned into magnetic domains, about 10 micrometers across,^[17] 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.^[16]

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

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 3.1×10²² years has been established.^[18]

⁶⁰Fe is an extinct radionuclide of long half-life (2.6 million years).^[19] It is not found on Earth, but its ultimate decay product is its granddaughter, the stable nuclide ⁶⁰Ni.^[18] 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.^[20]

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.^[21]

The most abundant iron isotope ⁵⁶Fe is of particular interest to nuclear scientists because it represents the most common endpoint of nucleosynthesis.^[22] 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, since addition of another alpha particle, resulting in ⁶⁰Zn, requires a great deal more energy. 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.^[23] It is also very common in the universe, relative to other stable metals of approximately the same atomic weight.^[23][24] Iron is the sixth most abundant element in the universe, and the most common refractory element.^[25]

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.^[26] Hence, elements heavier than iron require a supernova for their formation, involving rapid neutron capture by starting ⁵⁶Fe nuclei.^[23]

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.^[27]

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.^[28][29]

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.^[30] About 1 in 20 meteorites consist of the unique iron-nickel minerals taenite (35–80% iron) and kamacite (90–95% iron).^[31] 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.^[32]

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.^[33] 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.^[34] Silicate perovskite may form up to 93% of the lower mantle,^[35] 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.^[36]

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.^[37][38] 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).^[39]

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.^[40][41] 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.^[42]

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.^[43][44]

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).^[45] Through Eisensandstein (a jurassic 'iron sandstone', e.g. from Donzdorf in Germany)^[46] and Bath stone in the UK, iron compounds are responsible for the yellowish color of many historical buildings and sculptures.^[47] The proverbial red color of the surface of Mars is derived from an iron oxide-rich regolith.^[48]

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.^[49] In fact, iron is so common that production generally focuses only on ores with very high quantities of it.^[50]

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).^[51]

Oceans

Ocean science demonstrated the role of the iron in the ancient seas in both marine biota and climate.^[52]

Binary compounds

Oxides and sulfides

Ferrous or iron(II) oxide, FeO

Ferric or iron(III) oxide Fe₂O₃

Ferrosoferric or iron(II,III) oxide Fe₃O₄

Iron forms various oxide and hydroxide compounds; the most common are iron(II,III) oxide (Fe₃O₄), and iron(III) oxide (Fe₂O₃). Iron(II) oxide also exists, though it is unstable at room temperature. Despite their names, they are actually all non-stoichiometric compounds whose compositions may vary.^[62] These oxides are the principal ores for the production of iron (see bloomery and blast furnace). They are also used in the production of ferrites, useful magnetic storage media in computers, and pigments. The best known sulfide is iron pyrite (FeS₂), also known as fool's gold owing to its golden luster.^[58] It is not an iron(IV) compound, but is actually an iron(II) polysulfide containing Fe²⁺ and S²⁻
₂ ions in a distorted sodium chloride structure.^[62]

Pourbaix diagram of iron

Halides

Hydrated iron(III) chloride (ferric chloride)

The binary ferrous and ferric halides are well-known. The ferrous halides typically arise from treating iron metal with the corresponding hydrohalic acid to give the corresponding hydrated salts.^[58]

Fe + 2 HX → FeX₂ + H₂ (X = F, Cl, Br, I)

Iron reacts with fluorine, chlorine, and bromine to give the corresponding ferric halides, ferric chloride being the most common.^[63]

2 Fe + 3 X₂ → 2 FeX₃ (X = F, Cl, Br)

Ferric iodide is an exception, being thermodynamically unstable due to the oxidizing power of Fe³⁺ and the high reducing power of I⁻:^[63]

2 I⁻ + 2 Fe³⁺ → I₂ + 2 Fe²⁺ (E⁰ = +0.23 V)

Ferric iodide, a black solid, is not stable in ordinary conditions, but can be prepared through the reaction of iron pentacarbonyl with iodine and carbon monoxide in the presence of hexane and light at the temperature of −20 °C, with oxygen and water excluded.^[63]Complexes of ferric iodide with some soft bases are known to be stable compounds.^[64][65]

Solution chemistry

Comparison of colors of solutions of ferrate (left) and permanganate (right)

The standard reduction potentials in acidic aqueous solution for some common iron ions are given below:^[7]

The red-purple tetrahedral ferrate(VI) anion is such a strong oxidizing agent that it oxidizes ammonia to nitrogen (N₂) and water to oxygen^[63]

4 FeO²⁻
₄ + 34 H
₂O → 4 [Fe(H₂O)₆]³⁺ + 20 OH⁻
+ 3 O₂

The pale-violet hexaquo complex [Fe(H₂O)₆]³⁺ is an acid such that above pH 0 it is fully hydrolyzed:^[66]

Blue-green iron(II) sulfate heptahydrate

As pH rises above 0 the above yellow hydrolyzed species form and as it rises above 2–3, reddish-brown hydrous iron(III) oxide precipitates out of solution. Although Fe³⁺ has a d⁵ configuration, its absorption spectrum is not like that of Mn²⁺ with its weak, spin-forbidden d–d bands, because Fe³⁺ has higher positive charge and is more polarizing, lowering the energy of its ligand-to-metal charge transfer absorptions. Thus, all the above complexes are rather strongly colored, with the single exception of the hexaquo ion – and even that has a spectrum dominated by charge transfer in the near ultraviolet region.^[66] On the other hand, the pale green iron(II) hexaquo ion [Fe(H₂O)₆]²⁺ does not undergo appreciable hydrolysis. Carbon dioxide is not evolved when carbonate anions are added, which instead results in white iron(II) carbonate being precipitated out. In excess carbon dioxide this forms the slightly soluble bicarbonate, which occurs commonly in groundwater, but it oxidises quickly in air to form iron(III) oxide that accounts for the brown deposits present in a sizeable number of streams.^[67]

Coordination compounds

Due to its electronic structure, iron has a very large coordination and organometallic chemistry.

The two enantiomorphs of the ferrioxalate ion

Many coordination compounds of iron are known. A typical six-coordinate anion is hexachloroferrate(III), [FeCl₆]³⁻, found in the mixed salt tetrakis(methylammonium) hexachloroferrate(III) chloride.^[68][69] Complexes with multiple bidentate ligands have geometric isomers. For example, the trans-chlorohydridobis(bis-1,2-(diphenylphosphino)ethane)iron(II) complex is used as a starting material for compounds with the Fe(dppe)₂ moiety.^[70][71] The ferrioxalate ion with three oxalate ligands (shown at right) displays helical chirality with its two non-superposable geometries labelled Λ (lambda) for the left-handed screw axis and Δ (delta) for the right-handed screw axis, in line with IUPAC conventions.^[66] Potassium ferrioxalate is used in chemical actinometry and along with its sodium salt undergoes photoreduction applied in old-style photographic processes. The dihydrate of iron(II) oxalate has a polymeric structure with co-planar oxalate ions bridging between iron centres with the water of crystallisation located forming the caps of each octahedron, as illustrated below.^[72]

Crystal structure of iron(II) oxalate dihydrate, showing iron (gray), oxygen (red), carbon (black), and hydrogen (white) atoms.

Blood-red positive thiocyanate test for iron(III)

Iron(III) complexes are quite similar to those of chromium(III) with the exception of iron(III)'s preference for O-donor instead of N-donor ligands. The latter tend to be rather more unstable than iron(II) complexes and often dissociate in water. Many Fe–O complexes show intense colors and are used as tests for phenols or enols. For example, in the ferric chloride test, used to determine the presence of phenols, iron(III) chloride reacts with a phenol to form a deep violet complex:^[66]

3 ArOH + FeCl₃ → Fe(OAr)₃ + 3 HCl (Ar = aryl)

Among the halide and pseudohalide complexes, fluoro complexes of iron(III) are the most stable, with the colorless [FeF₅(H₂O)]²⁻ being the most stable in aqueous solution. Chloro complexes are less stable and favor tetrahedral coordination as in [FeCl₄]⁻; [FeBr₄]⁻ and [FeI₄]⁻ are reduced easily to iron(II). Thiocyanate is a common test for the presence of iron(III) as it forms the blood-red [Fe(SCN)(H₂O)₅]²⁺. Like manganese(II), most iron(III) complexes are high-spin, the exceptions being those with ligands that are high in the spectrochemical series such as cyanide. An example of a low-spin iron(III) complex is [Fe(CN)₆]³⁻. Iron shows a great variety of electronic spin states, including every possible spin quantum number value for a d-block element from 0 (diamagnetic) to 5⁄2 (5 unpaired electrons). This value is always half the number of unpaired electrons. Complexes with zero to two unpaired electrons are considered low-spin and those with four or five are considered high-spin.^[62]

Iron(II) complexes are less stable than iron(III) complexes but the preference for O-donor ligands is less marked, so that for example [Fe(NH₃)₆]²⁺ is known while [Fe(NH₃)₆]³⁺ is not. They have a tendency to be oxidized to iron(III) but this can be moderated by low pH and the specific ligands used.^[67]

Organometallic compounds

Iron penta-
carbonyl

Organoiron chemistry is the study of organometallic compounds of iron, where carbon atoms are covalently bound to the metal atom. They are many and varied, including cyanide complexes, carbonyl complexes, sandwich and half-sandwich compounds.

Prussian blue

Prussian blue or "ferric ferrocyanide", Fe₄[Fe(CN)₆]₃, is an old and well-known iron-cyanide complex, extensively used as pigment and in several other applications. Its formation can be used as a simple wet chemistry test to distinguish between aqueous solutions of Fe²⁺ and Fe³⁺ as they react (respectively) with potassium ferricyanide and potassium ferrocyanide to form Prussian blue.^[58]

Another old example of an organoiron compound is iron pentacarbonyl, Fe(CO)₅, in which a neutral iron atom is bound to the carbon atoms of five carbon monoxide molecules. The compound can be used to make carbonyl iron powder, a highly reactive form of metallic iron. Thermolysis of iron pentacarbonyl gives triiron dodecacarbonyl, Fe₃(CO)₁₂, a complex with a cluster of three iron atoms at its core. Collman's reagent, disodium tetracarbonylferrate, is a useful reagent for organic chemistry; it contains iron in the −2 oxidation state. Cyclopentadienyliron dicarbonyl dimer contains iron in the rare +1 oxidation state.^[73]

Structural formula of ferrocene and a powdered sample

A landmark in this field was the discovery in 1951 of the remarkably stable sandwich compound ferrocene Fe(C₅H₅)₂, by Pauson and Kealy^[74] and independently by Miller and colleagues,^[75] whose surprising molecular structure was determined only a year later by Woodward and Wilkinson^[76] and Fischer.^[77] Ferrocene is still one of the most important tools and models in this class.^[78]

Iron-centered organometallic species are used as catalysts. The Knölker complex, for example, is a transfer hydrogenation catalyst for ketones.^[79]

Industrial uses

The iron compounds produced on the largest scale in industry are iron(II) sulfate (FeSO₄·7H₂O) and iron(III) chloride (FeCl₃). The former is one of the most readily available sources of iron(II), but is less stable to aerial oxidation than Mohr's salt ((NH₄)₂Fe(SO₄)₂·6H₂O). Iron(II) compounds tend to be oxidized to iron(III) compounds in the air.^[58]

Development of iron metallurgy

Iron is one of the elements undoubtedly known to the ancient world.^[80] It has been worked, or wrought, for millennia. However, iron artefacts of great age are much rarer than objects made of gold or silver due to the ease with which iron corrodes.^[81] The technology developed slowly, and even after the discovery of smelting it took many centuries for iron to replace bronze as the metal of choice for tools and weapons.

Meteoritic iron

Iron harpoon head from Greenland. The iron edge covers a narwhal tusk harpoon using meteorite iron from the Cape York meteorite, one of the largest iron meteorites known.

Beads made from meteoric iron in 3500 BC or earlier were found in Gerzeh, Egypt by G.A. Wainwright.^[82] The beads contain 7.5% nickel, which is a signature of meteoric origin since iron found in the Earth's crust generally has only minuscule nickel impurities.

Meteoric iron was highly regarded due to its origin in the heavens and was often used to forge weapons and tools.^[82] For example, a dagger made of meteoric iron was found in the tomb of Tutankhamun, containing similar proportions of iron, cobalt, and nickel to a meteorite discovered in the area, deposited by an ancient meteor shower.^[83][84][85] Items that were likely made of iron by Egyptians date from 3000 to 2500 BC.^[81]

Meteoritic iron is comparably soft and ductile and easily cold forged but may get brittle when heated because of the nickel content.^[86]

Wrought iron

Main article: Wrought iron

Further information: Ancient iron production

The symbol for Mars has been used since antiquity to represent iron.

The iron pillar of Delhi is an example of the iron extraction and processing methodologies of early India.

The first iron production started in the Middle Bronze Age, but it took several centuries before iron displaced bronze. Samples of smelted iron from Asmar, Mesopotamia and Tall Chagar Bazaar in northern Syria were made sometime between 3000 and 2700 BC.^[87] The Hittites established an empire in north-central Anatolia around 1600 BC. They appear to be the first to understand the production of iron from its ores and regard it highly in their society.^[88] The Hittites began to smelt iron between 1500 and 1200 BC and the practice spread to the rest of the Near East after their empire fell in 1180 BC.^[87] The subsequent period is called the Iron Age.

Artifacts of smelted iron are found in India dating from 1800 to 1200 BC,^[89] and in the Levant from about 1500 BC (suggesting smelting in Anatolia or the Caucasus).^[90][91] Alleged references (compare history of metallurgy in South Asia) to iron in the Indian Vedas have been used for claims of a very early usage of iron in India respectively to date the texts as such. The rigveda term ayas (metal) refers to copper, while iron which is called as śyāma ayas, literally "black copper", first is mentioned in the post-rigvedic Atharvaveda.^[92]

Some archaeological evidence suggests iron was smelted in Zimbabwe and southeast Africa as early as the eighth century BC.^[93] Iron working was introduced to Greece in the late 11th century BC, from which it spread quickly throughout Europe.^[94]

Iron sickle from Ancient Greece

The spread of ironworking in Central and Western Europe is associated with Celtic expansion. According to Pliny the Elder, iron use was common in the Roman era.^[82] In the lands of what is now considered China, iron appears approximately 700–500 BC.^[95] Iron smelting may have been introduced into China through Central Asia.^[96] The earliest evidence of the use of a blast furnace in China dates to the 1st century AD,^[97] and cupola furnaces were used as early as the Warring States period (403–221 BC).^[98] Usage of the blast and cupola furnace remained widespread during the Tang and Song dynasties.^[99]

During the Industrial Revolution in Britain, Henry Cort began refining iron from pig iron to wrought iron (or bar iron) using innovative production systems. In 1783 he patented the puddling process for refining iron ore. It was later improved by others, including Joseph Hall.^[100]

Cast iron

Main article: Cast iron

Cast iron was first produced in China during 5th century BC,^[101] but was hardly in Europe until the medieval period.^[102][103] The earliest cast iron artifacts were discovered by archaeologists in what is now modern Luhe County, Jiangsu in China. Cast iron was used in ancient China for warfare, agriculture, and architecture.^[104] 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.^[105]

Coalbrookdale by Night, 1801. Blast furnaces light the iron making town of Coalbrookdale.

Medieval blast furnaces were about 10 feet (3.0 m) tall and made of fireproof brick; forced air was usually provided by hand-operated bellows.^[103] Modern blast furnaces have grown much bigger, with hearths fourteen meters in diameter that allow them to produce thousands of tons of iron each day, but essentially operate in much the same way as they did during medieval times.^[105]

In 1709, Abraham Darby I established a coke-fired blast furnace to produce cast iron, replacing charcoal, although continuing to use blast furnaces. The ensuing availability of inexpensive iron was one of the factors leading to the Industrial Revolution. Toward the end of the 18th century, cast iron began to replace wrought iron for certain purposes, because it was cheaper. Carbon content in iron was not implicated as the reason for the differences in properties of wrought iron, cast iron, and steel until the 18th century.^[87]

Since iron was becoming cheaper and more plentiful, it also became a major structural material following the building of the innovative first iron bridge in 1778. This bridge still stands today as a monument to the role iron played in the Industrial Revolution. Following this, iron was used in rails, boats, ships, aqueducts, and buildings, as well as in iron cylinders in steam engines.^[105] Railways have been central to the formation of modernity and ideas of progress^[106] and various languages refer to railways as iron road (e.g. French chemin de fer, German Eisenbahn, Turkish demiryolu, Russian железная дорога, Chinese, Japanese, and Korean 鐵道, Vietnamese đường sắt).

Foundations of modern chemistry

In 1774, Antoine Lavoisier used the reaction of water steam with metallic iron inside an incandescent iron tube to produce hydrogen in his experiments leading to the demonstration of the conservation of mass, which was instrumental in changing chemistry from a qualitative science to a quantitative one.^[109]

Laboratory routes

For a few limited purposes when it is needed, pure iron is produced in the laboratory in small quantities by reducing the pure oxide or hydroxide with hydrogen, or forming iron pentacarbonyl and heating it to 250 °C so that it decomposes to form pure iron powder.^[42] Another method is electrolysis of ferrous chloride onto an iron cathode.^[112]

Main industrial route

Iron production 2009 (million tonnes)^[113]

Country	Iron ore	Pig iron	Direct iron	Steel
China	1,114.9	549.4		573.6
Australia	393.9	4.4		5.2
Brazil	305.0	25.1	0.011	26.5
Japan		66.9		87.5
India	257.4	38.2	23.4	63.5
Russia	92.1	43.9	4.7	60.0
Ukraine	65.8	25.7		29.9
South Korea	0.1	27.3		48.6
Germany	0.4	20.1	0.38	32.7
World	1,594.9	914.0	64.5	1,232.4

Nowadays, the industrial production of iron or steel consists of two main stages. In the first stage, iron ore is reduced with coke in a blast furnace, and the molten metal is separated from gross impurities such as silicate minerals. This stage yields an alloy – pig iron – that contains relatively large amounts of carbon. In the second stage, the amount of carbon in the pig iron is lowered by oxidation to yield wrought iron, steel, or cast iron.^[114] Other metals can be added at this stage to form alloy steels.

Direct iron reduction

Owing to environmental concerns, alternative methods of processing iron have been developed. "Direct iron reduction" reduces iron ore to a ferrous lump called "sponge" iron or "direct" iron that is suitable for steelmaking.^[103] Two main reactions comprise the direct reduction process:

Natural gas is partially oxidized (with heat and a catalyst):^[103]

${\displaystyle {\ce {2 CH4 + O2 -> 2 CO + 4 H2}}}$

Iron ore is then treated with these gases in a furnace, producing solid sponge iron:^[103]

${\displaystyle {\ce {Fe2O3 + CO + 2 H2 -> 2 Fe + CO2 + 2 H2O}}}$

Silica is removed by adding a limestone flux as described above.^[103]

Thermite process

Ignition of a mixture of aluminium powder and iron oxide yields metallic iron via the thermite reaction:

${\displaystyle {\ce {Fe2O3 + 2 Al -> 2 Fe + Al2O3}}}$

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.^[105]

Molten oxide electrolysis

Molten oxide electrolysis uses an alloy of chromium, iron and other metals that does not react with oxygen and a liquid iron cathode while the electrolyte is a mixture of molten metal oxides into which iron ore is dissolved. The current keeps the electrolyte molten, and reduces the iron oxide. In addition to pure liquid iron, put oxygen is also produced, which can be sold to offset part of the cost. Production cell size is variable and can be much smaller than conventional furnaces. The only cardon dioxide emissions come from the electricity used to heat and reduce the metal.^[118]

As structural material

Iron is the most widely used of all the metals, accounting for over 90% of worldwide metal production. Its low cost and high strength often make it the material of choice to withstand stress or transmit forces, such as the construction of machinery and machine tools, rails, automobiles, ship hulls, concrete reinforcing bars, and the load-carrying framework of buildings. Since pure iron is quite soft, it is most commonly combined with alloying elements to make steel.^[121]

Types of steels and alloys

See also: Steel

Iron-carbon phase diagram

α-Iron is a fairly soft metal that can dissolve only a small concentration of carbon (no more than 0.021% by mass at 910 °C).^[125] Austenite (γ-iron) 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.^[17]

Commercially available iron is classified based on purity and the abundance of additives. Pig iron has 3.5–4.5% carbon^[126] and contains varying amounts of contaminants such as sulfur, silicon and phosphorus. Pig iron is not a saleable product, but rather an intermediate step in the production of cast iron and steel. The reduction of contaminants in pig iron that negatively affect material properties, such as sulfur and phosphorus, yields cast iron containing 2–4% carbon, 1–6% silicon, and small amounts of manganese.^[114] Pig iron 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.^[7] Its mechanical properties vary greatly and depend on the form the carbon takes in the alloy.^[13]

"White" cast irons contain their carbon in the form of cementite, or iron carbide (Fe₃C).^[13] 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 iron carbide, a very pale, silvery, shiny material, hence the appellation. Cooling a mixture of iron with 0.8% carbon slowly below 723 °C to room temperature results in separate, alternating layers of cementite and α-iron, which is soft and malleable and is called pearlite for its appearance. Rapid cooling, on the other hand, does not allow time for this separation and creates hard and brittle martensite. The steel can then be tempered by reheating to a temperature in between, changing the proportions of pearlite and martensite. The end product below 0.8% carbon content is a pearlite-αFe mixture, and that above 0.8% carbon content is a pearlite-cementite mixture.^[13]

In gray iron the carbon exists as separate, fine flakes of graphite, and also renders the material brittle due to the sharp edged flakes of graphite that produce stress concentration sites within the material.^[127] A newer variant of gray iron, referred to as ductile iron, is specially treated with trace amounts of magnesium to alter the shape of graphite to spheroids, or nodules, reducing the stress concentrations and vastly increasing the toughness and strength of the material.^[127]

Wrought iron contains less than 0.25% carbon but large amounts of slag that give it a fibrous characteristic.^[126] 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 within 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 than wrought iron, but is cheaper and more widely available. Carbon steel contains 2.0% carbon or less,^[128] 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 are usually only employed for specialist uses. One common alloy steel, though, is stainless steel. 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.^{[128][129][130]}

Alloys with high purity elemental makeups (such as alloys of electrolytic iron) have specifically enhanced properties such as ductility, tensile strength, toughness, fatigue strength, heat resistance, and corrosion resistance.

Apart from traditional applications, iron is also used for protection from ionizing radiation. Although it is lighter than another traditional protection material, lead, it is much stronger mechanically. The attenuation of radiation as a function of energy is shown in the graph.^[131]

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, a cost amounting to over 1% of the world's economy.^[132] Painting, galvanization, passivation, plastic coating and bluing are all used to protect iron from rust by excluding water and oxygen or by cathodic protection. The mechanism of the rusting of iron is as follows:^[132]

Cathode: 3 O₂ + 6 H₂O + 12 e⁻ → 12 OH⁻

Anode: 4 Fe → 4 Fe²⁺ + 8 e⁻; 4 Fe²⁺ → 4 Fe³⁺ + 4 e⁻

Overall: 4 Fe + 3 O₂ + 6 H₂O → 4 Fe³⁺ + 12 OH⁻ → 4 Fe(OH)₃ or 4 FeO(OH) + 4 H₂O

The electrolyte is usually iron(II) sulfate in urban areas (formed when atmospheric sulfur dioxide attacks iron), and salt particles in the atmosphere in seaside areas.^[132]

Catalysts and reagents

Because Fe is inexpensive and nontoxic, much effort has been devoted to the development of Fe-based catalysts and reagents. Iron is however less common as a catalyst in commercial processes than more expensive metals.^[133] In biology, Fe-containing enzymes are pervasive.^[134]

Iron catalysts are traditionally used in the Haber–Bosch process for the production of ammonia and the Fischer–Tropsch process for conversion of carbon monoxide to hydrocarbons for fuels and lubricants.^[135] Powdered iron in an acidic medium is used in the Bechamp reduction, the conversion of nitrobenzene to aniline.^[136]

Iron compounds

Iron(III) oxide mixed with aluminium powder can be ignited to create a thermite reaction, used in welding large iron parts (like rails) and purifying ores. Iron(III) oxide and oxyhydroxide are used as reddish and ocher pigments.

Iron(III) chloride finds use 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 circuit boards.^[137] It can also be dissolved in alcohol to form tincture of iron, which is used as a medicine to stop bleeding in canaries.^[138]

Iron(II) sulfate is used as a precursor to other iron compounds. It is also used to reduce chromate in cement. It is used to fortify foods and treat iron deficiency anemia. Iron(III) sulfate is used in settling minute sewage particles in tank water. Iron(II) chloride is used as a reducing flocculating agent, in the formation of iron complexes and magnetic iron oxides, and as a reducing agent in organic synthesis.^[137]

Sodium nitroprusside is a drug used as a vasodilator. It is on the World Health Organization's List of Essential Medicines.^[139]

Biochemistry

Iron acquisition poses a problem for aerobic organisms because ferric iron is poorly soluble near neutral pH. Thus, these organisms have developed means to absorb iron as complexes, sometimes taking up ferrous iron before oxidising it back to ferric iron.^[6] In particular, bacteria have evolved very high-affinity sequestering agents called siderophores.^{[146][147][148]}

After uptake in human cells, iron storage is precisely regulated.^[6][149] A major component of this regulation is the protein transferrin, which binds iron ions absorbed from the duodenum and carries it in the blood to cells.^[6][150] Transferrin contains Fe³⁺ in the middle of a distorted octahedron, bonded to one nitrogen, three oxygens and a chelating carbonate anion that traps the Fe³⁺ ion: it has such a high stability constant that it is very effective at taking up Fe³⁺ ions even from the most stable complexes. At the bone marrow, transferrin is reduced from Fe³⁺ and Fe²⁺ and stored as ferritin to be incorporated into hemoglobin.^[142]

The most commonly known and studied bioinorganic iron compounds (biological iron molecules) are the heme proteins: examples are hemoglobin, myoglobin, and cytochrome P450.^[6] These compounds participate in transporting gases, building enzymes, and transferring electrons.^[142] Metalloproteins are a group of proteins with metal ion cofactors. Some examples of iron metalloproteins are ferritin and rubredoxin.^[142] Many enzymes vital to life contain iron, such as catalase,^[151] lipoxygenases,^[152] and IRE-BP.^[153]

Hemoglobin is an oxygen carrier that occurs in red blood cells and contributes their color, transporting oxygen in the arteries from the lungs to the muscles where it is transferred to myoglobin, which stores it until it is needed for the metabolic oxidation of glucose, generating energy.^[6] Here the hemoglobin binds to carbon dioxide, produced when glucose is oxidized, which is transported through the veins by hemoglobin (predominantly as bicarbonate anions) back to the lungs where it is exhaled.^[142] In hemoglobin, the iron is in one of four heme groups and has six possible coordination sites; four are occupied by nitrogen atoms in a porphyrin ring, the fifth by an imidazole nitrogen in a histidine residue of one of the protein chains attached to the heme group, and the sixth is reserved for the oxygen molecule it can reversibly bind to.^[142] When hemoglobin is not attached to oxygen (and is then called deoxyhemoglobin), the Fe²⁺ ion at the center of the heme group (in the hydrophobic protein interior) is in a high-spin configuration. It is thus too large to fit inside the porphyrin ring, which bends instead into a dome with the Fe²⁺ ion about 55 picometers above it. In this configuration, the sixth coordination site reserved for the oxygen is blocked by another histidine residue.^[142]

When deoxyhemoglobin picks up an oxygen molecule, this histidine residue moves away and returns once the oxygen is securely attached to form a hydrogen bond with it. This results in the Fe²⁺ ion switching to a low-spin configuration, resulting in a 20% decrease in ionic radius so that now it can fit into the porphyrin ring, which becomes planar.^[142] (Additionally, this hydrogen bonding results in the tilting of the oxygen molecule, resulting in a Fe–O–O bond angle of around 120° that avoids the formation of Fe–O–Fe or Fe–O₂–Fe bridges that would lead to electron transfer, the oxidation of Fe²⁺ to Fe³⁺, and the destruction of hemoglobin.) This results in a movement of all the protein chains that leads to the other subunits of hemoglobin changing shape to a form with larger oxygen affinity. Thus, when deoxyhemoglobin takes up oxygen, its affinity for more oxygen increases, and vice versa.^[142] Myoglobin, on the other hand, contains only one heme group and hence this cooperative effect cannot occur. Thus, while hemoglobin is almost saturated with oxygen in the high partial pressures of oxygen found in the lungs, its affinity for oxygen is much lower than that of myoglobin, which oxygenates even at low partial pressures of oxygen found in muscle tissue.^[142] As described by the Bohr effect (named after Christian Bohr, the father of Niels Bohr), the oxygen affinity of hemoglobin diminishes in the presence of carbon dioxide.^[142]

A heme unit of human carboxyhemoglobin, showing the carbonyl ligand at the apical position, trans to the histidine residue^[154]

Carbon monoxide and phosphorus trifluoride are poisonous to humans because they bind to hemoglobin similarly to oxygen, but with much more strength, so that oxygen can no longer be transported throughout the body. Hemoglobin bound to carbon monoxide is known as carboxyhemoglobin. This effect also plays a minor role in the toxicity of cyanide, but there the major effect is by far its interference with the proper functioning of the electron transport protein cytochrome a.^[142] The cytochrome proteins also involve heme groups and are involved in the metabolic oxidation of glucose by oxygen. The sixth coordination site is then occupied by either another imidazole nitrogen or a methionine sulfur, so that these proteins are largely inert to oxygen – with the exception of cytochrome a, which bonds directly to oxygen and thus is very easily poisoned by cyanide.^[142] Here, the electron transfer takes place as the iron remains in low spin but changes between the +2 and +3 oxidation states. Since the reduction potential of each step is slightly greater than the previous one, the energy is released step-by-step and can thus be stored in adenosine triphosphate. Cytochrome a is slightly distinct, as it occurs at the mitochondrial membrane, binds directly to oxygen, and transports protons as well as electrons, as follows:^[142]

4 Cytc²⁺ + O₂ + 8H⁺
_inside → 4 Cytc³⁺ + 2 H₂O + 4H⁺
_outside

Although the heme proteins are the most important class of iron-containing proteins, the iron–sulfur proteins are also very important, being involved in electron transfer, which is possible since iron can exist stably in either the +2 or +3 oxidation states. These have one, two, four, or eight iron atoms that are each approximately tetrahedrally coordinated to four sulfur atoms; because of this tetrahedral coordination, they always have high-spin iron. The simplest of such compounds is rubredoxin, which has only one iron atom coordinated to four sulfur atoms from cysteine residues in the surrounding peptide chains. Another important class of iron–sulfur proteins is the ferredoxins, which have multiple iron atoms. Transferrin does not belong to either of these classes.^[142]

The ability of sea mussels to maintain their grip on rocks in the ocean is facilitated by their use of organometallic iron-based bonds in their protein-rich cuticles. Based on synthetic replicas, the presence of iron in these structures increased elastic modulus 770 times, tensile strength 58 times, and toughness 92 times. The amount of stress required to permanently damage them increased 76 times.^[155]

Nutrition

Deficiency

Iron deficiency is the most common nutritional deficiency in the world.^{[6][167][168][169]} When loss of iron is not adequately compensated by adequate dietary iron intake, a state of latent iron deficiency occurs, which over time leads to iron-deficiency anemia if left untreated, which is characterised by an insufficient number of red blood cells and an insufficient amount of hemoglobin.^[170] Children, pre-menopausal women (women of child-bearing age), and people with poor diet are most susceptible to the disease. Most cases of iron-deficiency anemia are mild, but if not treated can cause problems like fast or irregular heartbeat, complications during pregnancy, and delayed growth in infants and children.^[171]

The brain is resistant to acute iron deficiency due to the slow transport of iron through the blood brain barrier.^[172] Acute fluctuations in iron status (marked by serum ferritin levels) do not reflect brain iron status, but prolonged nutritional iron deficiency is suspected to reduce brain iron concentrations over time.^[173][174] In the brain, iron plays a role in oxygen transport, myelin synthesis, mitochondrial respiration, and as a cofactor for neurotransmitter synthesis and metabolism.^[175] Animal models of nutritional iron deficiency report biomolecular changes resembling those seen in Parkinson's and Huntington's disease.^[176][177] However, age-related accumulation of iron in the brain has also been linked to the development of Parkinson's.^[178]

Excess

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 primarily accomplished by regulating uptake.^[179] 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 and leads to abnormally low levels of hepcidin, a key regulator of the entry of iron into the circulatory system in mammals.^[180] In these people, excessive iron intake can result in iron overload disorders, known medically as hemochromatosis.^[6] Many people have an undiagnosed genetic susceptibility to iron overload, and are not aware of a family history of the problem. For this reason, people should not take iron supplements unless they suffer from iron deficiency and have consulted a doctor. Hemochromatosis is estimated to be the cause of 0.3–0.8% of all metabolic diseases of Caucasians.^[181]

Overdoses of ingested iron can cause excessive levels of free iron in the blood. High blood levels of free ferrous iron react with peroxides to produce highly reactive free radicals that can damage DNA, proteins, lipids, and other cellular components. Iron toxicity occurs when the cell contains free iron, which generally occurs when iron levels exceed the availability of transferrin to bind the iron. Damage to the cells of the gastrointestinal tract can also prevent them from regulating iron absorption, leading to further increases in blood levels. Iron typically damages cells in the heart, liver and elsewhere, causing adverse effects that include coma, metabolic acidosis, shock, liver failure, coagulopathy, long-term organ damage, and even death.^[182] Humans experience iron toxicity when the iron exceeds 20 milligrams for every kilogram of body mass; 60 milligrams per kilogram is considered a lethal dose.^[183] Overconsumption 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.^[183] The Dietary Reference Intake (DRI) sets the Tolerable Upper Intake Level (UL) for adults at 45 mg/day. For children under fourteen years old the UL is 40 mg/day.^[184]

The medical management of iron toxicity is complicated, and can include use of a specific chelating agent called deferoxamine to bind and expel excess iron from the body.^{[182][185][186]}

ADHD

Some research has suggested that low thalamic iron levels may play a role in the pathophysiology of ADHD.^[187] Some researchers have found that iron supplementation can be effective especially in the inattentive subtype of the disorder.^[188] One study also showed that iron may be able to decrease the risk of cardiovascular events during treatment with ADHD drugs.^[189]

Some researchers in the 2000s suggested a link between low levels of iron in the blood and ADHD. A 2012 study found no such correlation.^[190]

Cancer

The role of iron in cancer defense can be described as a "double-edged sword" because of its pervasive presence in non-pathological processes.^[191] People having chemotherapy may develop iron deficiency and anemia, for which intravenous iron therapy is used to restore iron levels.^[192] Iron overload, which may occur from high consumption of red meat,^[6] may initiate tumor growth and increase susceptibility to cancer onset,^[192] particularly for colorectal cancer.^[6]

Marine systems

Iron plays an essential role in marine systems and can act as a limiting nutrient for planktonic activity.^[193] Because of this, too much of a decrease in iron may lead to a decrease in growth rates in phytoplanktonic organisms such as diatoms.^[194] Iron can also be oxidized by marine microbes under conditions that are high in iron and low in oxygen.^[195]

Iron can enter marine systems through adjoining rivers and directly from the atmosphere. Once iron enters the ocean, it can be distributed throughout the water column through ocean mixing and through recycling on the cellular level.^[196] In the arctic, sea ice plays a major role in the store and distribution of iron in the ocean, depleting oceanic iron as it freezes in the winter and releasing it back into the water when thawing occurs in the summer.^[197] The iron cycle can fluctuate the forms of iron from aqueous to particle forms altering the availability of iron to primary producers.^[198] Increased light and warmth increases the amount of iron that is in forms that are usable by primary producers.^[199]