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This article is about the metallic element. For other uses, see Iron (disambiguation).
Iron,  26Fe
A rough wedge of silvery metal
Iron Spectrum.jpg
General properties
Name, symbol iron, Fe
Pronunciation /ˈ.ərn/
Appearance lustrous metallic with a grayish tinge
Iron in the periodic table
Hydrogen (diatomic nonmetal)
Helium (noble gas)
Lithium (alkali metal)
Beryllium (alkaline earth metal)
Boron (metalloid)
Carbon (polyatomic nonmetal)
Nitrogen (diatomic nonmetal)
Oxygen (diatomic nonmetal)
Fluorine (diatomic nonmetal)
Neon (noble gas)
Sodium (alkali metal)
Magnesium (alkaline earth metal)
Aluminium (post-transition metal)
Silicon (metalloid)
Phosphorus (polyatomic nonmetal)
Sulfur (polyatomic nonmetal)
Chlorine (diatomic nonmetal)
Argon (noble gas)
Potassium (alkali metal)
Calcium (alkaline earth metal)
Scandium (transition metal)
Titanium (transition metal)
Vanadium (transition metal)
Chromium (transition metal)
Manganese (transition metal)
Iron (transition metal)
Cobalt (transition metal)
Nickel (transition metal)
Copper (transition metal)
Zinc (transition metal)
Gallium (post-transition metal)
Germanium (metalloid)
Arsenic (metalloid)
Selenium (polyatomic nonmetal)
Bromine (diatomic nonmetal)
Krypton (noble gas)
Rubidium (alkali metal)
Strontium (alkaline earth metal)
Yttrium (transition metal)
Zirconium (transition metal)
Niobium (transition metal)
Molybdenum (transition metal)
Technetium (transition metal)
Ruthenium (transition metal)
Rhodium (transition metal)
Palladium (transition metal)
Silver (transition metal)
Cadmium (transition metal)
Indium (post-transition metal)
Tin (post-transition metal)
Antimony (metalloid)
Tellurium (metalloid)
Iodine (diatomic nonmetal)
Xenon (noble gas)
Caesium (alkali metal)
Barium (alkaline earth metal)
Lanthanum (lanthanide)
Cerium (lanthanide)
Praseodymium (lanthanide)
Neodymium (lanthanide)
Promethium (lanthanide)
Samarium (lanthanide)
Europium (lanthanide)
Gadolinium (lanthanide)
Terbium (lanthanide)
Dysprosium (lanthanide)
Holmium (lanthanide)
Erbium (lanthanide)
Thulium (lanthanide)
Ytterbium (lanthanide)
Lutetium (lanthanide)
Hafnium (transition metal)
Tantalum (transition metal)
Tungsten (transition metal)
Rhenium (transition metal)
Osmium (transition metal)
Iridium (transition metal)
Platinum (transition metal)
Gold (transition metal)
Mercury (transition metal)
Thallium (post-transition metal)
Lead (post-transition metal)
Bismuth (post-transition metal)
Polonium (post-transition metal)
Astatine (metalloid)
Radon (noble gas)
Francium (alkali metal)
Radium (alkaline earth metal)
Actinium (actinide)
Thorium (actinide)
Protactinium (actinide)
Uranium (actinide)
Neptunium (actinide)
Plutonium (actinide)
Americium (actinide)
Curium (actinide)
Berkelium (actinide)
Californium (actinide)
Einsteinium (actinide)
Fermium (actinide)
Mendelevium (actinide)
Nobelium (actinide)
Lawrencium (actinide)
Rutherfordium (transition metal)
Dubnium (transition metal)
Seaborgium (transition metal)
Bohrium (transition metal)
Hassium (transition metal)
Meitnerium (unknown chemical properties)
Darmstadtium (unknown chemical properties)
Roentgenium (unknown chemical properties)
Copernicium (transition metal)
Ununtrium (unknown chemical properties)
Flerovium (post-transition metal)
Ununpentium (unknown chemical properties)
Livermorium (unknown chemical properties)
Ununseptium (unknown chemical properties)
Ununoctium (unknown chemical properties)


Atomic number (Z) 26
Group, block group 8, d-block
Period period 4
Element category   transition metal
Standard atomic weight (±) (Ar) 55.845(2)[1]
Electron configuration [Ar] 3d6 4s2
per shell
2, 8, 14, 2
Physical properties
Phase solid
Melting point 1811 K ​(1538 °C, ​2800 °F)
Boiling point 3134 K ​(2862 °C, ​5182 °F)
Density near r.t. 7.874 g/cm3
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, +1,[2] +2, +3, +4, +5,[3] +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
Atomic radius empirical: 126 pm
Covalent radius Low spin: 132±3 pm
High spin: 152±6 pm
Crystal structure body-centered cubic (bcc)
Body-centered cubic crystal structure for iron

a=286.65 pm
Crystal structure face-centered cubic (fcc)
Face-centered cubic crystal structure for iron

between 1185–1667 K
Speed of sound thin rod 5120 m/s (at r.t.) (electrolytic)
Thermal expansion 11.8 µm/(m·K) (at 25 °C)
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
Poisson ratio 0.29
Mohs hardness 4
Vickers hardness 608 MPa
Brinell hardness 200–1180 MPa
CAS Number 7439-89-6
Discovery before 5000 BC
Most stable isotopes of iron
iso NA half-life DM DE (MeV) DP
54Fe 5.85% >3.1×1022 y εε 0.668 54Cr
55Fe syn 2.73 y ε 0.231 55Mn
56Fe 91.75% 56Fe is stable with 30 neutrons
57Fe 2.12% 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 2.6×106 y β 3.978 60Co
Decay modes in parentheses are predicted, but have not yet been observed
| references

Iron is a chemical element with symbol Fe (from Latin: ferrum, ultimately from ferre to bear or carry) and atomic number 26. It is a metal in the first transition series.[4] It is by mass the most common element on Earth, forming much of Earth's outer and inner core. It is the fourth most common element in the Earth's crust. Its abundance in rocky planets like Earth is due to its abundant production by fusion in high-mass stars, where the production of nickel-56 (which decays to the most common isotope of iron) is the last nuclear fusion reaction that is exothermic. Consequently, radioactive nickel is the last element to be produced before the violent collapse of a supernova, which scatters this precursor radionuclide of stable iron into space.

Like the other group 8 elements, ruthenium and osmium, iron exists in a wide range of oxidation states, −2 to +6, although +2 and +3 are the most common. Elemental iron occurs in meteoroids and other low oxygen environments, but is reactive to oxygen and water. Fresh iron surfaces appear lustrous silvery-gray, but oxidize in normal air to give hydrated iron oxides, commonly known as rust. Unlike the metals that form passivating oxide layers, iron oxides occupy more volume than the metal and thus flake off, exposing fresh surfaces for corrosion.

Iron metal has been used since ancient times, although copper alloys, which have lower melting temperatures, were used even earlier in human history. Pure iron is relatively soft, but is unobtainable by smelting. The material is significantly hardened and strengthened by impurities, in particular carbon, from the smelting process. A certain proportion of carbon (between 0.002% and 2.1%) produces steel, which may be up to 1000 times harder than pure iron. Crude iron metal is produced in blast furnaces, where ore is reduced by coke to pig iron, which has a high carbon content. Further refinement with oxygen reduces the carbon content to the correct proportion to make steel. Steels and iron alloys formed with other metals (alloy steels) are by far the most common industrial metals because they have a great range of desirable properties and iron-bearing rock is abundant.

Iron chemical compounds have many uses. Iron oxide mixed with aluminium powder can be ignited to create a thermite reaction, used in welding and purifying ores. Iron forms binary compounds with the halogens and the chalcogens. Among its organometallic compounds is ferrocene, the first sandwich compound discovered.

Iron plays an important role in biology, forming complexes with molecular oxygen in hemoglobin and myoglobin; these two compounds are common oxygen transport proteins in vertebrates. 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. A human male of average height has about 4 grams of iron in his body, a female about 3.5 grams. This iron is distributed throughout the body in hemoglobin, tissues, muscles, bone marrow, blood proteins, enzymes, ferritin, hemosiderin, and transport in plasma.[5]


Mechanical properties

Characteristic values of tensile strength (TS) and Brinell hardness (BH) of different forms of iron.[6][7]
Material TS
Iron whiskers 11000
Ausformed (hardened)
2930 850–1200
Martensitic steel 2070 600
Bainitic steel 1380 400
Pearlitic steel 1200 350
Cold-worked iron 690 200
Small-grain iron 340 100
Carbon-containing iron 140 40
Pure, single-crystal iron 10 3

The mechanical properties of iron and its alloys can be evaluated using a variety of tests, including the Brinell test, Rockwell test and the Vickers hardness test. The data on iron is so consistent that it is often used to calibrate measurements or to compare tests.[7][8] However, the mechanical properties of iron are significantly affected by the sample's purity: pure, single crystals of iron are actually softer than aluminium,[6] and the purest industrially produced iron (99.99%) has a hardness of 20–30 Brinell.[9] An increase in the carbon content will cause a significant increase in the iron's hardness and tensile strength. Maximum hardness of 65 Rc is achieved with a 0.6% carbon content, although the alloy has low tensile strength.[10] Because of the softness of iron, it is much easier to work with than its heavier congeners ruthenium and osmium.[11]

Molar volume vs. pressure for α iron at room temperature

Because of its significance for planetary cores, the physical properties of iron at high pressures and temperatures have also been studied extensively. The form of iron that is stable under standard conditions can be subjected to pressures up to ca. 15 GPa before transforming into a high-pressure form, as described in the next section.

Phase diagram and allotropes

Main article: Allotropes of iron

Iron represents an example of allotropy in a metal. There are at least four allotropic forms of iron, known as α, γ, δ, and ε; at very high pressures, some controversial experimental evidence exists for a phase β stable at very high pressures and temperatures.[12]

Low-pressure phase diagram of pure iron

As molten iron cools it crystallizes at 1538 °C 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, or ferrite. Finally, at 770 °C (the Curie point, Tc) 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 have the same axis orientation; however, the electrons of neighboring domains have other orientations with the result of mutual cancellation and no magnetic field. In magnetized iron, the electronic spins of the domains are aligned and the magnetic effects are reinforced. Although each domain contains billions of atoms, they are very small, about 10 micrometres across.[13] At pressures above approximately 10 GPa and temperatures of a few hundred kelvin or less, α-iron changes into a 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. The β-phase, if it exists, would appear at pressures of at least 50 GPa and temperatures of at least 1500 K and have an orthorhombic or a double hcp structure.[12]

Iron is of greatest importance when mixed with certain other metals and with carbon to form steels. There are many types of steel, all with different properties, and an understanding of the properties of the allotropes of iron is key to the manufacture of good quality steels.[14]

α-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).[15]

Above 912 °C and up to 1400 °C α-iron undergoes a phase transition from bcc to the fcc configuration of γ-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.[13]

The high-pressure phases of iron are important as endmember models for the solid parts of planetary cores. The inner core of the Earth is generally presumed to be an iron-nickel alloy with ε (or β) structure.[16]

The melting point of iron is experimentally well defined for pressures less than 50 GPa. For greater pressures, studies put the γ-ε-liquid triple point at pressures that differ by tens of gigapascals and 1000 K in the melting point. Generally speaking, molecular dynamics computer simulations of iron melting and shock wave experiments suggest higher melting points and a much steeper slope of the melting curve than static experiments carried out in diamond anvil cells.[17] 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; however, they are higher than the values for the previous element manganese because that element has a half-filled 3d subshell and consequently its d-electrons are not easily delocalized.[11]


Main article: Isotopes of iron

Naturally occurring iron consists of four stable isotopes: 5.845% of 54Fe, 91.754% of 56Fe, 2.119% of 57Fe and 0.282% of 58Fe. Of these stable isotopes, only 57Fe has a nuclear spin (−12). The nuclide 54Fe theoretically can undergo double beta decay, 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 nickel-60.[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. It is often cited, falsely, as the isotope of highest binding energy, a distinction which actually belongs to nickel-62.[22] Since 56Ni is 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 particle 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 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 cobalt-56, and then to stable iron-56. This last nuclide is therefore common in the universe, relative to other stable metals of approximately the same atomic weight.[23][24]

Nuclei of iron atoms have some of the highest binding energies per nucleon, surpassed only by the nickel isotope 62Ni. It 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. Element production in supernovas and distribution on Earth greatly favor iron over nickel.[25]

Iron-56 is the heaviest stable isotope produced by the alpha process in stellar nucleosynthesis; elements heavier than iron and nickel 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.[24]


Iron is created by extremely large stars with extremely hot (over 2.5 billion kelvin) cores through the silicon burning process. It is the heaviest stable element to be produced in this manner. The process starts with the second largest stable nucleus created by silicon burning, which is calcium. One stable nucleus of calcium fuses with one helium nucleus, creating unstable titanium. Before the titanium decays, it can fuse with another helium nucleus, creating unstable chromium. Before the chromium decays, it can fuse with another helium nucleus, creating unstable iron. Before the iron decays, it can fuse with another helium nucleus, creating unstable nickel-56. Any further fusion of nickel-56 consumes energy instead of producing energy, so after the production of nickel-56, the star does not produce the energy necessary to keep the core from collapsing. Eventually, the nickel-56 decays to unstable cobalt-56, which in turn decays to stable iron-56. When the core of the star collapses, it creates a supernova. Supernovas also create additional stable iron isotopes via the r-process.[23]


Planetary occurrence

Iron meteorites, similar in composition to the Earth's inner- and outer core

Iron is the sixth most abundant element in the Universe, and the most common refractory element.[26] It is formed as the final exothermic stage of stellar nucleosynthesis, by silicon fusion in massive stars.[23]

Metallic or native iron is rarely found on the surface of the Earth because it tends to oxidize, but its oxides are pervasive and represent the primary ores. While it makes up about 5% of the Earth's crust, both the Earth's inner and outer core are 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.[27][28] Most of the iron in the crust is found combined with oxygen as iron oxide minerals such as hematite (Fe2O3) and magnetite (Fe3O4). Large deposits of iron are found in banded iron formations. These geological formations are 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[29][30]

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 main form of natural metallic iron on the Earth's surface.[31]

The red color of the surface of Mars is derived from an iron oxide-rich regolith. This has been proven by Mössbauer spectroscopy.[32]

Stocks in use in society

According to the International Resource Panel's Metal Stocks in Society report, the global stock of iron in use in society is 2200 kg per capita. Much of this is in more-developed countries (7000–14000 kg per capita) rather than less-developed countries (2000 kg per capita).[33]

Chemistry and compounds

Representative compound
−2 Disodium tetracarbonylferrate (Collman's reagent)
−1 Fe
0 Iron pentacarbonyl
1 Cyclopentadienyliron dicarbonyl dimer ("Fp2")
2 Ferrous sulfate, ferrocene
3 Ferric chloride, ferrocenium tetrafluoroborate
4 Fe(diars)
5 FeO3−
6 Potassium ferrate

Iron forms compounds mainly in the +2 and +3 oxidation states. Traditionally, iron(II) compounds are called ferrous, and iron(III) compounds ferric. Iron also occurs in higher oxidation states, an example being the purple potassium ferrate (K2FeO4) which contains iron in its +6 oxidation state, although this is very easily reduced. Iron(IV) is a common intermediate in many biochemical oxidation reactions.[34][35] Numerous organometallic compounds contain formal oxidation states of +1, 0, −1, or even −2. The oxidation states and other bonding properties are often assessed using the technique of Mössbauer spectroscopy.[36] There are also many mixed valence compounds that contain both iron(II) and iron(III) centers, such as magnetite and Prussian blue (Fe4(Fe[CN]6)3).[35] The latter is used as the traditional "blue" in blueprints.[37]

Iron is the first of the transition metals that cannot reach its group oxidation state of +8, although its heavier congeners ruthenium and osmium can, with ruthenium having more difficulty than osmium.[38] While iron's most common oxidation states are +2 and +3, ruthenium's is +3 and osmium's is +4. Iron also commonly forms aqueous cations in the +2 and +3 oxidation states, which is possible for ruthenium but not osmium.[38]

Some canary-yellow powder sits, mostly in lumps, on a laboratory watch glass.
Hydrated iron(III) chloride, also known as ferric chloride

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.[35]

Unlike many other metals, iron does not form amalgams with mercury. As a result, mercury is traded in standardized 76 pound flasks (34 kg) made of iron.[39]

Iron is by far the most reactive element in its group; it is pyrophoric when finely divided and dissolves easily in dilute acids, giving Fe2+. However, it does not react with concentrated nitric acid and other oxidizing acids due to the formation of an impervious oxide layer, which can nevertheless react with hydrochloric acid.[38]

The standard reduction potentials in acidic aqueous solution for some common iron ions are given below:[38]

Fe2+ + 2e ⇌ Fe E0 = −0.447 V
Fe3+ + 3e ⇌ Fe E0 = −0.037 V
+ 8H+ + 3e
⇌ Fe3+ + 4H2O E0 = +2.20 V

The red-purple ferrate(VI) anion is such a strong oxidizing agent that it oxidizes nitrogen and ammonia at room temperature, and even water itself in acidic or neutral solutions:[40]

4 FeO2−
+ 10 H
→ 4 Fe3+
+ 20 OH
+ 3 O2

Binary compounds

Iron reacts with oxygen in the air to form 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. 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.[35]

The binary ferrous and ferric halides are well-known, with the exception of ferric iodide. The ferrous halides typically arise from treating iron metal with the corresponding hydrohalic acid to give the corresponding hydrated salts.[35]

Fe + 2 HX → FeX2 + H2

Iron reacts with fluorine, chlorine, and bromine to give the corresponding ferric halides, ferric chloride being the most common.[40]

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.[40]

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

Coordination compounds

The two enantiomorphs of the ferrioxalate ion

Many coordination compounds of iron are known. A typical six-coordinate anions is hexachloroferrate(III), [FeCl6]3−, found in the mixed salt tetrakis(methylammonium) hexachloroferrate(III) chloride.[41][42] 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.[43][44] 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.[45] 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.[46]

Ball-and-stick model of a chain in the crystal structure of iron(II) oxalate dihydrate

Prussian blue, Fe4[Fe(CN)6]3 is the most famous of the cyanide complexes of iron. 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.[35] It can be used as an antidote for thallium and radioactive caesium poisoning.[47][48] Prussian blue can be used in laundry bluing to correct the yellowish tint left by ferrous salts in water.

Organometallic compounds

Fulvalene, which Pauson and Kealy sought to prepare
The (incorrect) structure for ferrocene that Pauson and Kealy proposed
The structural formula of ferrocene

Cyanide complexes are technically organometallic but more important are carbonyl complexes and sandwich and half-sandwich compounds. The premier iron(0) compound is iron pentacarbonyl, Fe(CO)5, which is used to produce carbonyl iron powder, a highly reactive form of metallic iron. Thermolysis of iron pentacarbonyl gives the trinuclear cluster, triiron dodecacarbonyl. 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.[49]

Ferrocene was first synthesised in 1951 during an attempt to prepare the fulvalene (C10H8) by oxidative dimerization of cyclopentadiene; the resultant product was found to have molecular formula C10H10Fe and reported to exhibit "remarkable stability".[50] The discovery sparked substantial interest in the field of organometallic chemistry,[51][52] in part because the structure proposed by Pauson and Kealy (shown at right) was inconsistent with then-existing bonding models and did not explain its unexpected stability. Consequently, the initial challenge was to definitively determine the structure of ferrocene in the hope that its bonding and properties would then be understood. The shockingly novel sandwich structure, [Fe(η5-C5H5)2],[51] was deduced and reported independently by three groups in 1952: Robert Burns Woodward and Geoffrey Wilkinson investigated the reactivity in order to determine the structure[53] and demonstrated that ferrocene undergoes similar reactions to a typical aromatic molecule (such as benzene),[54] Ernst Otto Fischer deduced the sandwich structure and also began synthesising other metallocenes including cobaltocene;[55] Eiland and Pepinsky provided X-ray crystallographic confirmation of the sandwich structure.[56] Applying valence bond theory to ferrocene by considering an Fe2+ centre and two cyclopentadienide anions (C5H5), which are known to be aromatic according to Hückel's rule and hence highly stable, allowed correct prediction of the geometry of the molecule. Once molecular orbital theory was successfully applied and the Dewar-Chatt-Duncanson model proposed,[57] the reasons for ferrocene's remarkable stability became clear.[58] Ferrocene was not the first organometallic compound known – Zeise's salt, K[PtCl3(C2H4)]·H2O was reported in 1831[59][60] and Mond's discovery of Ni(CO)4 occurred in 1888[61]   but it was ferrocene's discovery that began organometallic chemistry as a separate area of chemistry. It was so important that Wilkinson and Fischer shared the 1973 Nobel Prize for Chemistry "for their pioneering work, performed independently, on the chemistry of the organometallic, so called sandwich compounds".[62] Ferrocene itself can be used as the backbone of a ligand, e.g. dppf. Ferrocene can itself be oxidized to the ferrocenium cation (Fc+); the ferrocene/ferrocenium couple is often used as a reference in electrochemistry.[52]

Metallocenes like ferrocene can be prepared by reaction of Freshly-cracked cyclopentadiene with iron(II) chloride and a weak base.[63] It is an aromatic substance and undergoes substitution reactions rather than addition reactions on the cyclopentadienyl ligands. For example, Friedel-Crafts acylation of ferrocene with acetic anhydride yields acetylferrocene[64] just as acylation of benzene yields acetophenone under similar conditions.

Synthesis of acetylferrocene from dicyclopentadiene.png

Iron-centred organometallic species are used as catalysts. The Knölker complex, for example, is a transfer hydrogenation catalyst for ketones.[65]


Wrought iron

Further information: Ancient iron production
A circle, with a short, simple arrow shape extending diagonally upwards and rightwards from its edge
The symbol for Mars has been used since antiquity to represent iron.
An pillar, slightly fluted, with some ornamentation at its top. It is black, slightly weathered to a dark brown near the base. It is around 7 meters (23 feet) tall. It stands upon a raised circular base of stone, and is surrounded by a short, square fence.
The Iron pillar of Delhi is an example of the iron extraction and processing methodologies of early India.

Iron has been worked, or wrought, for millennia. However, iron objects of great age are much rarer than objects made of gold or silver due to the ease of corrosion of iron.[66] Beads made from meteoric iron in 3500 BCE or earlier were found in Gerzah, Egypt by G. A. Wainwright.[67] The beads contain 7.5% nickel, which is a signature of meteoric origin since iron found in the Earth's crust has very little to no nickel content. Meteoric iron was highly regarded due to its origin in the heavens and was often used to forge weapons and tools or whole specimens placed in churches.[67] Items that were likely made of iron by Egyptians date from 2500 to 3000 BCE.[66] Iron had a distinct advantage over bronze in warfare implements. It was much harder and more durable than bronze, although susceptible to rust. However, this is contested. Hittitologist Trevor Bryce argues that before advanced iron-working techniques were developed in India, meteoritic iron weapons used by early Mesopotamian armies had a tendency to shatter in combat, due to their high carbon content.[68]

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 2700 and 3000 BCE.[69] The Hittites appear to be the first to understand the production of iron from its ores and regard it highly in their society. They began to smelt iron between 1500 and 1200 BCE and the practice spread to the rest of the Near East after their empire fell in 1180 BCE.[69] The subsequent period is called the Iron Age. Iron smelting, and thus the Iron Age, reached Europe two hundred years later and arrived in Zimbabwe, Africa by the 8th century.[69] In China, iron only appears circa 700–500 BCE.[70] Iron smelting may have been introduced into China through Central Asia.[71] The earliest evidence of the use of a blast furnace in China dates to the 1st century AD,[72] and cupola furnaces were used as early as the Warring States period (403–221 BCE).[73] Usage of the blast and cupola furnace remained widespread during the Song and Tang Dynasties.[74]

Artifacts of smelted iron are found in India dating from 1800 to 1200 BCE,[75] and in the Levant from about 1500 BCE (suggesting smelting in Anatolia or the Caucasus).[76][77]

The Book of Genesis, fourth chapter, verse 22 contains the first mention of iron in the Old Testament of the Bible; "Tubal-cain, an instructor of every artificer in brass and iron."[66] Other verses allude to iron mining (Job 28:2), iron used as a stylus (Job 19:24), furnace (Deuteronomy 4:20), chariots (Joshua 17:16), nails (I Chron. 22:3), saws and axes (II Sam. 12:31), and cooking utensils (Ezekiel 4:3).[78] The metal is also mentioned in the New Testament, for example in Acts chapter 12 verse 10, "[Peter passed through] the iron gate that leadeth unto the city" of Antioch.[79]

Iron working was introduced to Greece in the late 11th century BCE.[80] 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.[67] The annual iron output of the Roman Empire is estimated at 84,750 t,[81] while the similarly populous Han China produced around 5,000 t.[82]

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.[83]

Cast iron

Cast iron was first produced in China during 5th century BCE,[84] but was hardly in Europe until the medieval period.[85][86] 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.[87] 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.

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.[86] Modern blast furnaces have grown much bigger.

In 1709, Abraham Darby I established a coke-fired blast furnace to produce cast iron. 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.[69]

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.[88]


See also: Steelmaking

Steel (with smaller carbon content than pig iron but more than wrought iron) was first produced in antiquity by using a bloomery. Blacksmiths in Luristan in western Iran were making good steel by 1000 BCE.[69] Then improved versions, Wootz steel by India and Damascus steel were developed around 300 BCE and 500 CE respectively. These methods were specialized, and so steel did not become a major commodity until the 1850s.[89]

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 made steel much more economical, thereby leading to wrought iron no longer being produced.[90]

Foundations of modern chemistry

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 mass conservation. Anaerobic oxidation of iron at high temperature can be schematically represented by the following reactions:

Fe + H2O → FeO + H2
2 Fe + 3 H2O → Fe2O3 + 3 H2
3 Fe + 4 H2O → Fe3O4 + 4 H2

Production of metallic iron

Industrial routes

See also: Iron ore

The production of iron or steel is a process consisting of two main stages, unless the desired product is cast iron. In the first stage pig iron is produced in a blast furnace. Alternatively, it may be directly reduced. In the second stage, pig iron is converted to wrought iron or steel.[91]

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.
How iron was extracted in the 19th century

For a few limited purposes when it is needed, pure iron is produced 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.[14]

Blast furnace processing

Main article: Blast furnace

Industrial iron production starts with iron ores, principally hematite, which has a nominal formula Fe2O3, and magnetite, with the formula Fe3O4. These ores are reduced to the metal in a carbothermic reaction, i.e. by treatment with carbon. The conversion is typically conducted in a blast furnace at temperatures of about 2000 °C. Carbon is provided in the form of coke. The process also contains a flux such as limestone, which is used to remove silicaceous minerals in the ore, which would otherwise clog the furnace. The coke and limestone are fed into the top of the furnace, while a massive blast of heated air, about 4 tons per ton of iron,[86] is forced into the furnace at the bottom.[91]

In the furnace, the coke reacts with oxygen in the air blast to produce carbon monoxide:[91]

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:[91]

Fe2O3 + 3 CO → 2 Fe + 3 CO2

Some iron in the high-temperature lower region of the furnace reacts directly with the coke:[91]

2 Fe2O3 + 3 C → 4 Fe + 3 CO2

The flux present to melt impurities in the ore is principally limestone (calcium carbonate) and dolomite (calcium-magnesium carbonate). Other specialized fluxes are used depending on the details of the ore. In the heat of the furnace the limestone flux decomposes to calcium oxide (also known as quicklime):[91]

CaCO3 → CaO + CO2

Then calcium oxide combines with silicon dioxide to form a liquid slag.[91]

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.[86]

This heap of iron ore pellets will be used in steel production.

Direct iron reduction

Owing to environmental concerns, alternative methods of processing iron have been developed. "Direct iron reduction" reduces iron ore to a powder called "sponge" iron or "direct" iron that is suitable for steelmaking.[86] Two main reactions comprise the direct reduction process:

Natural gas is partially oxidized (with heat and a catalyst):[86]

2 CH4 + O2 → 2 CO + 4 H2

These gases are then treated with iron ore in a furnace, producing solid sponge iron:[86]

Fe2O3 + CO + 2 H2 → 2 Fe + CO2 + 2 H2O

Silica is removed by adding a limestone flux as described above.[86]

Further processes

Main articles: Steelmaking and Ironworks
Iron-carbon phase diagram, various stable solid solution forms

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.[91] This form of iron, also known as cast 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.[88]

Annealing involves the heating of a piece of steel to 700–800 °C for several hours and then gradual cooling. It makes the steel softer and more workable.[92]

Laboratory methods

Metallic iron is generally produced in the laboratory by two methods. One route is electrolysis of ferrous chloride onto an iron cathode. The second method involves reduction of iron oxides with hydrogen gas at about 500 °C.[93]



See also: Steel
Iron production 2009 (million tonnes)[94]
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

Iron is the most widely used of all the metals, accounting for over 90% 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 combined with alloying elements to make steel.[95]

Commercially available iron is classified based on purity and the abundance of additives. Pig iron has 3.5–4.5% carbon[96] 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.[91] 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 and depend on the form the carbon takes in the alloy.

"White" cast irons contain their carbon in the form of cementite, or iron carbide (Fe3C).[11] 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.

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.[97] 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.[97]

Wrought iron contains less than 0.25% carbon but large amounts of slag that give it a fibrous characteristic.[96] 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,[98] 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.

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.

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. Painting, galvanization, passivation, plastic coating and bluing are all used to protect iron from rust by excluding water and oxygen or by cathodic protection.[99]

Iron compounds

Although its metallurgical role is dominant in terms of amounts, iron compounds are pervasive in industry as well being used in many niche uses. 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.[100] Powdered iron in an acidic solvent was used in the Bechamp reduction the reduction of nitrobenzene to aniline.[101]

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.[102] It can also be dissolved in alcohol to form tincture of iron. The other halides tend to be limited to laboratory uses.

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. These are its main uses. 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.

Biological role

Iron is involved in numerous biological processes.[103][104] Iron-proteins are found in all living organisms: archaeans, bacteria and eukaryotes, including humans. For example, the color of blood is due to the hemoglobin, an iron-containing protein. As illustrated by hemoglobin, iron is often bound to cofactors, e.g. in hemes. The iron-sulfur clusters are pervasive and include nitrogenase, the enzymes responsible for biological nitrogen fixation. Influential theories of evolution have invoked a role for iron sulfides in the iron-sulfur world theory.[105]

Structure of Heme b; in the protein additional ligand(s) would be attached to Fe.

Iron is a necessary trace element found in nearly all living organisms. Iron-containing enzymes and proteins, often containing heme prosthetic groups, participate in many biological oxidations and in transport. Examples of proteins found in higher organisms include hemoglobin, cytochrome (see high-valent iron), and catalase.[106]

Bioinorganic compounds

The most commonly known and studied bioinorganic iron compounds (biological iron molecules) are the heme proteins: examples are hemoglobin, myoglobin, and cytochrome P450. These compounds participate in transporting gases, building enzymes, and transferring electrons.[105] Metalloproteins are a group of proteins with metal ion cofactors. Some examples of iron metalloproteins are ferritin and rubredoxin.[105] Many enzymes vital to life contain iron, such as catalase, lipoxygenases, and IRE-BP.

Health and diet

Iron is pervasive, but particularly rich sources of dietary iron include red meat, lentils, beans, poultry, fish, leaf vegetables, watercress, tofu, chickpeas, black-eyed peas, blackstrap molasses, fortified bread, and fortified breakfast cereals. Iron in low amounts is found in molasses, teff, and farina. Iron in meat (heme iron) is more easily absorbed than iron in vegetables.[107] Although some studies suggest that heme/hemoglobin from red meat has effects which may increase the likelihood of colorectal cancer,[108][109] there is still some controversy[110] with a few studies suggesting that such claims are not supported by sufficient evidence.[111]

Iron provided by dietary supplements is often found as iron(II) fumarate, although iron sulfate is cheaper and is absorbed equally well. Elemental iron, or reduced iron, despite being absorbed at only one third to two thirds the efficiency (relative to iron sulfate),[112] is often added to foods such as breakfast cereals or enriched wheat flour. Iron is most available to the body when chelated to amino acids[113] and is also available for use as a common iron supplement. Glycine, the cheapest and most common amino acid is most often used to produce iron glycinate supplements.[114] The Recommended Dietary Allowance (RDA) for iron varies considerably depending on age, sex, and source of dietary iron (heme-based iron has higher bioavailability).[115] Infants may require iron supplements if they are bottle-fed cow's milk.[116] Blood donors and pregnant women are at special risk of low iron levels and are often advised to supplement their iron intake.[117]

Uptake and storage

Iron acquisition poses a problem for aerobic organisms because ferric iron is poorly soluble near neutral pH. Thus, bacteria have evolved high-affinity sequestering agents called siderophores.[118][119][120]

After uptake in cells, iron storage is carefully regulated; iron ions are never "free". 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.[121] In animals, plants, and fungi, iron is often 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.[105]

Inorganic iron 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 ions have a high potential for biological toxicity.[122]

Regulation of uptake

Main article: Hepcidin

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.[123] 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, known medically as hemochromatosis. 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 a cause of 0.3 to 0.8% of all metabolic diseases of Caucasians.[124]

MRIfstudies show that iron accumulates in the hippocampus of the brains of those with Alzheimer's disease and in the substantia nigra of those with Parkinson disease.[125]


Iron-eating bacteria live in the hulls of sunken ships such as the Titanic.[126] The acidophile bacteria Acidithiobacillus ferrooxidans, Leptospirillum ferrooxidans, Sulfolobus spp., Acidianus brierleyi and Sulfobacillus thermosulfidooxidans can oxidize ferrous iron enzymically.[127] A sample of the fungus Aspergillus niger was found growing from gold mining solution, and was found to contain cyano metal complexes such as gold, silver, copper iron and zinc. The fungus also plays a role in the solubilization of heavy metal sulfides.[128]

Permeable reactive barriers

Zerovalent iron is the main reactive material for permeable reactive barriers.[129]


NFPA 704
"fire diamond"
Flammability code 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g., canola oil Health code 0: Exposure under fire conditions would offer no hazard beyond that of ordinary combustible material. E.g., sodium chloride Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium Special hazards (white): no codeNFPA 704 four-colored diamond
Fire diamond for powdered iron metal
Main article: Iron poisoning

Overdoses of ingested iron can cause excessive levels of 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, adult respiratory distress syndrome, long-term organ damage, and even death.[130] 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.[131] 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.[131] 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.[132]

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.[130][133][134]

See also


  1. ^ Standard Atomic Weights 2013. Commission on Isotopic Abundances and Atomic Weights
  2. ^ Ram, R. S. & Bernath, P. F. (2003). "Fourier transform emission spectroscopy of the g4Δ-a4Δ system of FeCl" (PDF). Journal of Molecular Spectroscopy 221 (2): 261. Bibcode:2003JMoSp.221..261R. doi:10.1016/S0022-2852(03)00225-X. 
  3. ^ Demazeau, G.; Buffat, B.; Pouchard, M.; Hagenmuller, P. (1982). "Recent developments in the field of high oxidation states of transition elements in oxides stabilization of Six-coordinated Iron(V)". Zeitschrift für anorganische und allgemeine Chemie 491: 60. doi:10.1002/zaac.19824910109. 
  4. ^
  5. ^ "Iron in human health". 
  6. ^ a b Kohl, Walter H. (1995). Handbook of materials and techniques for vacuum devices. Springer. pp. 164–167. ISBN 1-56396-387-6. 
  7. ^ a b Kuhn, Howard and Medlin, Dana (prepared under the direction of the ASM International Handbook Committee), eds. (2000). ASM Handbook – Mechanical Testing and Evaluation (PDF) 8. ASM International. p. 275. ISBN 0-87170-389-0. 
  8. ^ "Hardness Conversion Chart". Maryland Metrics. Retrieved 23 May 2010. 
  9. ^ Takaji, Kusakawa; Toshikatsu, Otani (1964). "Properties of Various Pure Irons: Study on pure iron I". Tetsu-to-Hagane 50 (1): 42–47. 
  10. ^ Raghavan, V. (2004). Materials Science and Engineering. PHI Learning Pvt. Ltd. p. 218. ISBN 81-203-2455-2. 
  11. ^ a b c Greenwood and Earnshaw, pp. 1074–5
  12. ^ a b Boehler, Reinhard (2000). "High-pressure experiments and the phase diagram of lower mantle and core materials". Review of Geophysics (American Geophysical Union) 38 (2): 221–245. Bibcode:2000RvGeo..38..221B. doi:10.1029/1998RG000053. 
  13. ^ a b Bramfitt, B. L.; Benscoter, Arlan O. (2002). "The Iron Carbon Phase Diagram". Metallographer's guide: practice and procedures for irons and steels. ASM International. pp. 24–28. ISBN 978-0-87170-748-2. 
  14. ^ a b Greenwood and Earnshaw, pp. 1071–2
  15. ^ Martin, John Wilson (2007). Concise encyclopedia of the structure of materials. Elsevier. p. 183. ISBN 0-08-045127-6. 
  16. ^ Stixrude, Lars; Wasserman, Evgeny; Cohen, Ronald E. (1997-11-10). "Composition and temperature of Earth's inner core". Journal of Geophysical Research: Solid Earth 102 (B11): 24729–24739. doi:10.1029/97JB02125. ISSN 2156-2202. 
  17. ^ Boehler, Reinhard; Ross, M. (2007). "Properties of Rocks and Minerals_High-Pressure Melting". Mineral Physics. Treatise on Geophysics 2. Elsevier. pp. 527–541. doi:10.1016/B978-044452748-6.00047-X. 
  18. ^ a b Audi, G.; Bersillon, O.; Blachot, J.; Wapstra, A.H. (2003). "The NUBASE evaluation of nuclear and decay properties" (PDF). Nuclear Physics A 729: 3–128. Bibcode:2003NuPhA.729....3A. doi:10.1016/j.nuclphysa.2003.11.001. 
  19. ^ Rugel, G.; Faestermann, T.; Knie, K.; Korschinek, G.; Poutivtsev, M.; Schumann, D.; Kivel, N.; Günther-Leopold, I.; Weinreich, R.; Wohlmuther, M. (2009). "New Measurement of the 60Fe Half-Life". Physical Review Letters 103 (7). doi:10.1103/PhysRevLett.103.072502. ISSN 0031-9007. 
  20. ^ Dauphas, N.; Rouxel, O. (2006). "Mass spectrometry and natural variations of iron isotopes" (PDF). Mass Spectrometry Reviews 25 (4): 515–550. doi:10.1002/mas.20078. PMID 16463281. 
  21. ^ Mostefaoui, S.; Lugmair, G.W.; Hoppe, P.; El Goresy, A. (2004). "Evidence for live 60Fe in meteorites". New Astronomy Reviews 48: 155. Bibcode:2004NewAR..48..155M. doi:10.1016/j.newar.2003.11.022. 
  22. ^ Fewell, M. P. (1995). "The atomic nuclide with the highest mean binding energy". American Journal of Physics 63 (7): 653. Bibcode:1995AmJPh..63..653F. doi:10.1119/1.17828. 
  23. ^ a b c Woosley, S.; Janka, T. (2006). "The physics of core collapse supernovae". arXiv:astro-ph/0601261. 
  24. ^ a b Greenwood and Earnshaw, p. 12
  25. ^ Bautista, Manuel A.; Pradhan, Anil K. (1995). "Iron and Nickel Abundances in H~II Regions and Supernova Remnants". Bulletin of the American Astronomical Society 27: 865. Bibcode:1995AAS...186.3707B. 
  26. ^ McDonald, I.; Sloan, G. C.; Zijlstra, A. A.; Matsunaga, N.; Matsuura, M.; Kraemer, K. E.; Bernard-Salas, J.; Markwick, A. J. (2010). "Rusty Old Stars: A Source of the Missing Interstellar Iron?". The Astrophysical Journal Letters 717 (2): L92–L97. arXiv:1005.3489. Bibcode:2010ApJ...717L..92M. doi:10.1088/2041-8205/717/2/L92. 
  27. ^ "Iron: geological information". WebElements. Retrieved 23 May 2010. 
  28. ^ John W. Morgan & Edward Anders (1980). "Chemical composition of Earth, Venus, and Mercury". Proc. Natl. Acad. Sci. 77 (12): 6973–6977. Bibcode:1980PNAS...77.6973M. doi:10.1073/pnas.77.12.6973. PMC 350422. PMID 16592930. 
  29. ^ Lyons, T. W.; Reinhard, CT (2009). "Early Earth: Oxygen for heavy-metal fans". Nature 461 (7261): 179–181. Bibcode:2009Natur.461..179L. doi:10.1038/461179a. PMID 19741692. 
  30. ^ Cloud, P. (1973). "Paleoecological Significance of the Banded Iron-Formation". Economic Geology 68 (7): 1135–1143. doi:10.2113/gsecongeo.68.7.1135. 
  31. ^ Emiliani, Cesare (1992). "Planet earth: cosmology, geology, and the evolution of life and environment". Cambridge University Press: 152. ISBN 978-0-521-40949-0.  |chapter= ignored (help)
  32. ^ Klingelhöfer, G.; Morris, R. V.; Souza, P. A.; Rodionov, D.; Schröder, C. (2007). "Two earth years of Mössbauer studies of the surface of Mars with MIMOS II". Hyperfine Interactions 170: 169–177. Bibcode:2006HyInt.170..169K. doi:10.1007/s10751-007-9508-5. 
  33. ^ Metal Stocks in Society: Scientific synthesis, 2010, International Resource Panel, UNEP
  34. ^ Nam, Wonwoo (2007). "High-Valent Iron(IV)–Oxo Complexes of Heme and Non-Heme Ligands in Oxygenation Reactions". Accounts of Chemical Research 40 (7): 522–531. doi:10.1021/ar700027f. PMID 17469792. 
  35. ^ a b c d e f Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils (1985). "Iron". Lehrbuch der Anorganischen Chemie (in German) (91–100 ed.). Walter de Gruyter. pp. 1125–1146. ISBN 3-11-007511-3. 
  36. ^ Reiff, William Michael; Long, Gary J. (1984). "Mössbauer Spectroscopy and the Coordination Chemistry of Iron". Mössbauer spectroscopy applied to inorganic chemistry. Springer. pp. 245–283. ISBN 978-0-306-41647-7. 
  37. ^ Ware, Mike (1999). "An introduction in monochrome". Cyanotype: the history, science and art of photographic printing in Prussian blue. NMSI Trading Ltd. pp. 11–19. ISBN 978-1-900747-07-3. 
  38. ^ a b c d Greenwood and Earnshaw, pp. 1075–9
  39. ^ Gmelin, Leopold (1852). "Mercury and Iron". Hand-book of chemistry 6. Cavendish Society. pp. 128–129. 
  40. ^ a b c Greenwood and Earnshaw, p. 1082–4
  41. ^ Clausen, C. A.; Good, M. L. (1968). "Stabilization of the hexachloroferrate(III) anion by the methylammonium cation". Inorganic Chemistry 7 (12): 2662–2663. doi:10.1021/ic50070a047. 
  42. ^ James, B. D.; Bakalova, M.; Lieseganga, J.; Reiff, W. M.; Hockless, D. C. R.; Skelton, B. W.; White, A. H. (1996). "The hexachloroferrate(III) anion stabilized in hydrogen bonded packing arrangements. A comparison of the X-ray crystal structures and low temperature magnetism of tetrakis(methylammonium) hexachloroferrate(III) chloride (I) and tetrakis(hexamethylenediammonium) hexachloroferrate(III) tetrachloroferrate(III) tetrachloride (II)". Inorganica Chimica Acta 247 (2): 169–174. doi:10.1016/0020-1693(95)04955-X. 
  43. ^ Giannoccaro, P.; Sacco, A. (1977). "Bis[ethylenebis(diphenylphosphine)]-Hydridoiron Complexes". Inorg. Synth. 17: 69. doi:10.1002/9780470132487.ch19. 
  44. ^ Lee, J.; Jung, G.; Lee, S. W. (1998). "Structure of trans-chlorohydridobis(diphenylphosphinoethane)iron(II)". Bull. Korean Chem. Soc. 19 (2): 267–269. 
  45. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9. 
  46. ^ Echigo, Takuya; Kimata, Mitsuyoshi (2008). "Single-crystal X-ray diffraction and spectroscopic studies on humboldtine and lindbergite: weak Jahn–Teller effect of Fe2+ ion". Phys. Chem. Minerals 35: 467–475. doi:10.1007/s00269-008-0241-7. 
  47. ^ "Questions and Answers on Prussian Blue". Retrieved 6 June 2009. 
  48. ^ Thompson, D. F; Callen, ED (2004). "Soluble or Insoluble Prussian Blue for Radiocesium and Thallium Poisoning?". Annals of Pharmacotherapy 38 (9): 1509–1514. doi:10.1345/aph.1E024. PMID 15252192. 
  49. ^ Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. pp. 1282–86. ISBN 0-08-022057-6. .
  50. ^ Kealy, T. J.; Pauson, P. L. (1951). "A New Type of Organo-Iron Compound". Nature 168 (4285): 1039–1040. Bibcode:1951Natur.168.1039K. doi:10.1038/1681039b0. 
  51. ^ a b Laszlo, P.; Hoffmann, R. (2000). "Ferrocene: Ironclad History of Rashomon Tale?" (PDF). Angew. Chem. Int. Ed. 39 (1): 123–124. doi:10.1002/(SICI)1521-3773(20000103)39:1<123::AID-ANIE123>3.0.CO;2-Z. PMID 10649350. 
  52. ^ a b Federman Neto, A.; Pelegrino, A. C.; Darin, V. A. (2004). "Ferrocene: 50 Years of Transition Metal Organometallic Chemistry — From Organic and Inorganic to Supramolecular Chemistry". ChemInform 35 (43). doi:10.1002/chin.200443242.  (Abstract; original published in Trends Organomet. Chem., 4:147–169, 2002)
  53. ^ Wilkinson, G.; Rosenblum, M.; Whiting, M. C.; Woodward, R. B. (1952). "The Structure of Iron Bis-Cyclopentadienyl". J. Am. Chem. Soc. 74 (8): 2125–2126. doi:10.1021/ja01128a527. 
  54. ^ Werner, H. (2008). Landmarks in Organo-Transition Metal Chemistry: A Personal View. New York: Springer Science. pp. 161–163. ISBN 978-0-387-09847-0. 
  55. ^ Fischer, E. O.; Pfab, W. (1952). "Zur Kristallstruktur der Di-Cyclopentadienyl-Verbindungen des zweiwertigen Eisens, Kobalts und Nickels". Z. Anorg. Allg. Chem. (in German) 7 (6): 377–379. doi:10.1002/zaac.19532740603. 
  56. ^ Eiland, P. F.; Pepinsky, R. (1952). "X-ray Examination of Iron Biscyclopentadienyl". J. Am. Chem. Soc. 74 (19): 4971. doi:10.1021/ja01139a527. 
  57. ^ Mingos, D. M. P. (2001). "A Historical Perspective on Dewar's Landmark Contribution to Organometallic Chemistry". J. Organomet. Chem. 635 (1–2): 1–8. doi:10.1016/S0022-328X(01)01155-X. 
  58. ^ Mehrotra, R. C.; Singh, A. (2007). Organometallic Chemistry: A Unified Approach (2nd ed.). New Delhi: New Age International. pp. 261–267. ISBN 978-81-224-1258-1. 
  59. ^ Zeise, W. C. (1831). "Von der Wirkung zwischen Platinchlorid und Alkohol, und von den dabei entstehenden neuen Substanzen". Annalen der Physik (in German) 97 (4): 497–541. Bibcode:1831AnP....97..497Z. doi:10.1002/andp.18310970402. 
  60. ^ Hunt, L. B. (1984). "The First Organometallic Compounds: William Christopher Zeise and his Platinum Complexes" (PDF). Platinum Metals Rev. 28 (2): 76–83. 
  61. ^ Leigh, G. J.; Winterton, N., eds. (2002). "Section D: Transition Metal Complexes of Olefins, Acetylenes, Arenes and Related Isolobal COmpounds". Modern Coordination Chemistry: The Legacy of Joseph Chatt. Cambridge, UK: RSC Publishing. pp. 101–110. ISBN 0-85404-469-8. 
  62. ^ "The Nobel Prize in Chemistry 1973". Nobel Foundation. Retrieved 12 September 2010. 
  63. ^ Wilkinson, G. (1956). "Ferrocene". Org. Synth. 36: 31. doi:10.15227/orgsyn.036.0031. 
  64. ^ Bozak, R. E. (1966). "Acetylation of Ferrocene: A Chromatography Experiment for Elementary Organic Laboratory". J. Chem. Educ. 43 (2): 73. doi:10.1021/ed043p73. 
  65. ^ Bullock, R. M. (11 September 2007). "An Iron Catalyst for Ketone Hydrogenations under Mild Conditions". Angew. Chem. Int. Ed. 46 (39): 7360–7363. doi:10.1002/anie.200703053. 
  66. ^ a b c Weeks 1968, p. 29.
  67. ^ a b c Weeks 1968, p. 31.
  68. ^ Bryce, Trevor (2007). Hittite Warrior. Osprey Publishing. pp. 22–23. ISBN 978-1-84603-081-9. 
  69. ^ a b c d e Weeks 1968, p. 32.
  70. ^ Sawyer, Ralph D. and Mei-chün Sawyer. The Seven Military Classics of Ancient China. Boulder: Westview, (1993), p. 10.
  71. ^ Pigott, Vincent C. (1999). p. 8.
  72. ^ Peter J. Golas (25 February 1999). Science and Civilisation in China: Volume 5, Chemistry and Chemical Technology, Part 13, Mining. Cambridge University Press. p. 152. ISBN 978-0-521-58000-7. earlist blast furnace discovered in China from about the first century AD 
  73. ^ Pigott, Vincent C. (1999). The Archaeometallurgy of the Asian Old World. Philadelphia: University of Pennsylvania Museum of Archaeology and Anthropology. ISBN 0-924171-34-0, p. 191.
  74. ^ The Coming of the Ages of Steel. Brill Archive. 1961. p. 54. GGKEY:DN6SZTCNQ3G. 
  75. ^ Tewari, Rakesh. "The origins of Iron Working in India: New evidence from the Central Ganga plain and the Eastern Vindhyas" (PDF). State Archaeological Department. Retrieved 23 May 2010. 
  76. ^ Photos, E. (1989). "The Question of Meteoritic versus Smelted Nickel-Rich Iron: Archaeological Evidence and Experimental Results". World Archaeology (Taylor & Francis, Ltd.) 20 (3): 403–421. doi:10.1080/00438243.1989.9980081. JSTOR 124562. 
  77. ^ Muhly, James D. (2003). "Metalworking/Mining in the Levant". In Lake, Richard Winona. Near Eastern Archaeology IN: Eisenbrauns 180. pp. 174–183. 
  78. ^ Weeks 1968, pp. 29–30.
  79. ^ Weeks 1968, p. 30.
  80. ^ Riederer, Josef; Wartke, Ralf-B.: "Iron", Cancik, Hubert; Schneider, Helmuth (eds.): Brill's New Pauly, Brill 2009
  81. ^ Craddock, Paul T. (2008): "Mining and Metallurgy", in: Oleson, John Peter (ed.): The Oxford Handbook of Engineering and Technology in the Classical World, Oxford University Press, ISBN 978-0-19-518731-1, p. 108
  82. ^ Wagner, Donald B.: "The State and the Iron Industry in Han China", NIAS Publishing, Copenhagen 2001, ISBN 87-87062-77-1, p. 73
  83. ^ R. A. Mott, 'Dry and Wet Puddling' Trans. Newcomen Soc. 49, (1977–8), 156–7.
  84. ^ Wagner, Donald B. (2003). "Chinese blast furnaces from the 10th to the 14th century". Historical Metallurgy 37 (1): 25–37.  originally published in Wagner, Donald B. (2001). "Chinese blast furnaces from the 10th to the 14th century". West Asian Science, Technology, and Medicine 18: 41–74. 
  85. ^ Giannichedda, Enrico (2007): "Metal production in Late Antiquity", in Technology in Transition AD 300–650 Lavan, L.; Zanini, E. and Sarantis, A.(eds.), Brill, Leiden; ISBN 90-04-16549-5, p. 200.
  86. ^ a b c d e f g h Biddle, Verne; Parker, Gregory. Chemistry, Precision and Design. A Beka Book, Inc. 
  87. ^ Donald B. Wagner (1993). Iron and Steel in Ancient China. BRILL. pp. 335–340. ISBN 978-90-04-09632-5. 
  88. ^ a b Greenwood and Earnshaw, p. 1072
  89. ^ Spoerl, Joseph S. A Brief History of Iron and Steel Production. Saint Anselm College
  90. ^ Enghag, Per (8 January 2008). Encyclopedia of the Elements: Technical Data - History - Processing - Applications. pp. 190–191. ISBN 978-3-527-61234-5. 
  91. ^ a b c d e f g h i Greenwood and Earnshaw, p. 1073
  92. ^ Verhoeven, J.D. Fundamentals of Physical Metallurgy, Wiley, New York, 1975, p. 326
  93. ^ H. Lux "Metallic Iron" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 2. p. 1490–1.
  94. ^ Steel Statistical Yearbook 2010. World Steel Association
  95. ^ Greenwood and Earnshaw, pp. 1070–1
  96. ^ a b Camp, James McIntyre; Francis, Charles Blaine (1920). The Making, Shaping and Treating of Steel. Pittsburgh: Carnegie Steel Company. pp. 173–174. ISBN 1-147-64423-3. 
  97. ^ a b Smith, William F.; Hashemi, Javad (2006), Foundations of Materials Science and Engineering (4th ed.), McGraw-Hill, p. 431, ISBN 0-07-295358-6. 
  98. ^ "Classification of Carbon and Low-Alloy Steels". Retrieved 5 January 2008. 
  99. ^ Greenwood and Earnshaw, p. 1076
  100. ^ Kolasinski, Kurt W. (2002). "Where are Heterogenous Reactions Important". Surface science: foundations of catalysis and nanoscience. John Wiley and Sons. pp. 15–16. ISBN 978-0-471-49244-3. 
  101. ^ McKetta, John J. (1989). "Nitrobenzene and Nitrotoluene". Encyclopedia of Chemical Processing and Design: Volume 31 – Natural Gas Liquids and Natural Gasoline to Offshore Process Piping: High Performance Alloys. CRC Press. pp. 166–167. ISBN 978-0-8247-2481-8. 
  102. ^ Wildermuth, Egon; Stark, Hans; Friedrich, Gabriele; Ebenhöch, Franz Ludwig; Kühborth, Brigitte; Silver, Jack; Rituper, Rafael (2000). "Ullmann's Encyclopedia of Industrial Chemistry". doi:10.1002/14356007.a14_591. ISBN 3-527-30673-0.  |chapter= ignored (help)
  103. ^ Dlouhy, Adrienne C.; Outten, Caryn E. (2013). "Chapter 8 The Iron Metallome in Eukaryotic Organisms". In Banci, Lucia. Metallomics and the Cell. Metal Ions in Life Sciences 12. Springer. doi:10.1007/978-94-007-5561-1_8. ISBN 978-94-007-5560-4.  electronic-book ISBN 978-94-007-5561-1 ISSN 1559-0836 electronic-ISSN 1868-0402
  104. ^ Yee, Gereon M.; Tolman, William B. (2015). "Chapter 5 Transition Metal Complexes and the Activation of Dioxygen". In Peter M.H. Kroneck; Martha E. Sosa Torres. Sustaining Life on Planet Earth: Metalloenzymes Mastering Dioxygen and Other Chewy Gases. Metal Ions in Life Sciences 15. Springer. pp. 131–204. doi:10.1007/978-3-319-12415-5_5. 
  105. ^ a b c d Greenwood and Earnshaw, pp. 1098–1104
  106. ^ Lippard, S. J.; Berg, J. M. (1994). Principles of Bioinorganic Chemistry. Mill Valley: University Science Books. ISBN 0-935702-73-3. 
  107. ^ Food Standards Agency – Eat well, be well – Iron deficiency. (5 March 2012). Retrieved on 27 June 2012.
  108. ^ Sesink, Aloys L. A.; T; K; V (1999). "Red meat and colon cancer: the cytotoxic and hyperproliferative effects of dietary heme". Cancer Research 59 (22): 5704–9. PMID 10582688. 
  109. ^ Glei, M.; Klenow, S.; Sauer, J.; Wegewitz, U.; Richter, K.; Pool-Zobel, B. L. (2006). "Hemoglobin and hemin induce DNA damage in human colon tumor cells HT29 clone 19A and in primary human colonocytes". Mutat. Res. 594 (1–2): 162–171. doi:10.1016/j.mrfmmm.2005.08.006. PMID 16226281. 
  110. ^ Sandhu, M. S.; White, I. R.; McPherson, K. (2001). "Systematic Review of the Prospective Cohort Studies on Meat Consumption and Colorectal Cancer Risk: A Meta-Analytical Approach". Cancer Epidemiology, Biomarkers & Prevention 10 (5): 439–46. PMID 11352852. 
  111. ^ "Eating Red Meat Will Not Increase Colorectal Cancer Risk, Study Suggests". ScienceDaily. 13 June 2007. Retrieved 23 May 2010. 
  112. ^ Hoppe, M.; Hulthén, L.; Hallberg, L. (2005). "The relative bioavailability in humans of elemental iron powders for use in food fortification". European Journal of Nutrition 45 (1): 37–44. doi:10.1007/s00394-005-0560-0. PMID 15864409. 
  113. ^ Pineda, O.; Ashmead, H. D. (2001). "Effectiveness of treatment of iron-deficiency anemia in infants and young children with ferrous bis-glycinate chelate". Nutrition 17 (5): 381–4. doi:10.1016/S0899-9007(01)00519-6. PMID 11377130. 
  114. ^ Ashmead, H. DeWayne (1989). Conversations on Chelation and Mineral Nutrition. Keats Publishing. ISBN 0-87983-501-X. 
  115. ^ "Dietary Reference Intakes: Elements" (PDF). The National Academies. 2001. Retrieved 21 May 2008. 
  116. ^ "Iron Deficiency Anemia". MediResource. Retrieved 17 December 2008. 
  117. ^ Milman, N (1996). "Serum ferritin in Danes: studies of iron status from infancy to old age, during blood donation and pregnancy". International Journal of Hematology 63 (2): 103–35. doi:10.1016/0925-5710(95)00426-2. PMID 8867722. 
  118. ^ Neilands, JB (1995). "Siderophores: structure and function of microbial iron transport compounds". The Journal of Biological Chemistry 270 (45): 26723–6. doi:10.1074/jbc.270.45.26723. PMID 7592901. 
  119. ^ Neilands, J B (1981). "Microbial Iron Compounds". Annual Review of Biochemistry 50 (1): 715–31. doi:10.1146/ PMID 6455965. 
  120. ^ Boukhalfa, Hakim; Crumbliss, Alvin L. (2002). "Chemical aspects of siderophore mediated iron transport". BioMetals 15 (4): 325–39. doi:10.1023/A:1020218608266. PMID 12405526. 
  121. ^ Rouault, Tracey A. (2003). "How Mammals Acquire and Distribute Iron Needed for Oxygen-Based Metabolism". PLoS Biology 1 (3): e9. doi:10.1371/journal.pbio.0000079. PMC 212690. PMID 14551907. 
  122. ^ Nanami, M.; Ookawara, T; Otaki, Y; Ito, K; Moriguchi, R; Miyagawa, K; Hasuike, Y; Izumi, M; Eguchi, H; Suzuki, K; Nakanishi, T (2005). "Tumor necrosis factor-α-induced iron sequestration and oxidative stress in human endothelial cells". Arteriosclerosis, thrombosis, and vascular biology 25 (12): 2495–2501. doi:10.1161/01.ATV.0000190610.63878.20. PMID 16224057. 
  123. ^ Ramzi S. Cotran; Vinay Kumar; Tucker Collins; Stanley Leonard Robbins (1999). Robbins pathologic basis of disease. Saunders. ISBN 978-0-7216-7335-6. Retrieved 27 June 2012. 
  124. ^ Durupt, S; Durieu, I; Nové-Josserand, R; Bencharif, L; Rousset, H; Vital Durand, D (2000). "Hereditary hemochromatosis". Rev Med Interne 21 (11): 961–71. doi:10.1016/S0248-8663(00)00252-6. PMID 11109593. 
  125. ^ Brar, S; Henderson, D; Schenck, J; Zimmerman, EA (2009). "Iron accumulation in the substantia nigra of patients with Alzheimer disease and parkinsonism". Archives of neurology 66 (3): 371–4. doi:10.1001/archneurol.2008.586. PMID 19273756. 
  126. ^ Ward, Greg (2012). The Rough Guide to the Titanic. London: Rough Guides Ltd. p. 171. ISBN 978-1-4053-8699-9. 
  127. ^ Geoffrey Michael Gadd (March 2010). "Metals, minerals and microbes: geomicrobiology and bioremediation". Microbiology 156 (3): 609–643. doi:10.1099/mic.0.037143-0. PMID 20019082. 
  128. ^ Harbhajan Singh. Mycoremediation: Fungal Bioremediation. p. 509. 
  129. ^ Roehl, K.E.; Meggyes, T; Simon, F.G.; Stewart, D.I. (27 April 2005). Long-Term Performance of Permeable Reactive Barriers. p. 5. ISBN 978-0-08-053561-6. 
  130. ^ a b Cheney, K.; Gumbiner, C.; Benson, B.; Tenenbein, M. (1995). "Survival after a severe iron poisoning treated with intermittent infusions of deferoxamine". J Toxicol Clin Toxicol 33 (1): 61–6. doi:10.3109/15563659509020217. PMID 7837315. 
  131. ^ a b "Toxicity, Iron". Medscape. Retrieved 23 May 2010. 
  132. ^ Dietary Reference Intakes (DRIs): Recommended Intakes for Individuals (PDF), Food and Nutrition Board, Institute of Medicine, National Academies, 2004, retrieved 2009-06-09 
  133. ^ Tenenbein, M (1996). "Benefits of parenteral deferoxamine for acute iron poisoning". J Toxicol Clin Toxicol 34 (5): 485–489. doi:10.3109/15563659609028005. PMID 8800185. 
  134. ^ Wu H, Wu T, Xu X, Wang J, Wang J (May 2011). "Iron toxicity in mice with collagenase-induced intracerebral hemorrhage". J Cereb Blood Flow Metab. 31 (5): 1243–50. doi:10.1038/jcbfm.2010.209. PMC 3099628. PMID 21102602. 


Further reading

  • 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.

External links