Extracting usable metal fromiron ores requireskilns orfurnaces capable of reaching 1,500 °C (2,730 °F), about 500 °C (932 °F) higher than that required tosmeltcopper. Humans started to master that process inEurasia during the2nd millennium BC and the use of irontools andweapons began to displacecopper alloys – in some regions, only around 1200 BC. That event is considered the transition from theBronze Age to theIron Age. In themodern world, iron alloys, such assteel,stainless steel,cast iron andspecial steels, are by far the most common industrial metals, due to their mechanical properties and low cost. Theiron and steel industry is thus very important economically, and iron is the cheapest metal, with a price of a few dollars per kilogram or pound.
Pristine and smooth pure iron surfaces are a mirror-like silvery-gray. Iron reacts readily with oxygen andwater to produce brown-to-blackhydratediron oxides, commonly known asrust. Unlike the oxides of some other metals that formpassivating layers, rust occupies more volume than the metal and thus flakes off, exposing more fresh surfaces for corrosion. Chemically, the most common oxidation states of iron areiron(II) andiron(III). Iron shares many properties of other transition metals, including the othergroup 8 elements,ruthenium andosmium. Iron forms compounds in a wide range ofoxidation states, −4 to +7. Iron also forms manycoordination complexes; some of them, such asferrocene,ferrioxalate, andPrussian blue have substantial industrial, medical, or research applications.
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 denotedα,γ,δ, andε.
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 abody-centered cubic (bcc)crystal structure. As it cools further to 1394 °C, it changes to its γ-iron allotrope, aface-centered cubic (fcc) crystal structure, oraustenite. At 912 °C and below, the crystal structure again becomes the bcc α-iron allotrope.[11]
The physical properties of iron at very high pressures and temperatures have also been studied extensively,[12][13] 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 anotherhexagonal close-packed (hcp) structure, which is also known asε-iron. The higher-temperature γ-phase also changes into ε-iron,[13] but does so at higher pressure.
Some controversial experimental evidence exists for a stableβ phase at pressures above 50 GPa and temperatures of at least 1500 K. It is supposed to have anorthorhombic or a double hcp structure.[14] (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.[11])
The melting and boiling points of iron, along with itsenthalpy of atomization, are lower than those of the earlier 3d elements fromscandium tochromium, 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;[16] however, they are higher than the values for the previous elementmanganese because that element has a half-filled 3d sub-shell and consequently its d-electrons are not easily delocalized. This same trend appears forruthenium but notosmium.[17]
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.[18]
Below itsCurie point of 770 °C (1,420 °F; 1,040 K), α-iron changes fromparamagnetic toferromagnetic: thespins of the two unpaired electrons in each atom generally align with the spins of its neighbors, creating an overallmagnetic field.[20] This happens because the orbitals of those two electrons (dz2 and dx2 −y2) do not point toward neighboring atoms in the lattice, and therefore are not involved in metallic bonding.[11]
In the absence of an external source of magnetic field, the atoms get spontaneously partitioned intomagnetic domains, about 10 micrometers across,[21] 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 aselectrical transformers,magnetic recording heads, andelectric 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.[20]
Similar behavior is exhibited by some iron compounds, such as theferrites including the mineralmagnetite, a crystalline form of the mixed iron(II,III) oxideFe3O4 (although the atomic-scale mechanism,ferrimagnetism, is somewhat different). Pieces of magnetite with natural permanent magnetization (lodestones) provided the earliestcompasses for navigation. Particles of magnetite were extensively used in magnetic recording media such ascore memories,magnetic tapes,floppies, anddisks, until they were replaced bycobalt-based materials.
Iron has four stableisotopes:54Fe (5.845% of natural iron),56Fe (91.754%),57Fe (2.119%) and58Fe (0.282%). Twenty-four artificial isotopes have also been created. Of these stable isotopes, only57Fe has anuclear spin (−1⁄2). Thenuclide54Fe theoretically can undergodouble electron capture to54Cr, but the process has never been observed and only a lower limit on the half-life of 4.4×1020 years has been established.[22]
60Fe is anextinct radionuclide of longhalf-life (2.6 million years).[23] It is not found on Earth, but its ultimate decay product is its granddaughter, the stable nuclide60Ni.[9] Much of the past work on isotopic composition of iron has focused on thenucleosynthesis of60Fe through studies ofmeteorites and ore formation. In the last decade, advances inmass spectrometry have allowed the detection and quantification of minute, naturally occurring variations in the ratios of thestable isotopes of iron. Much of this work is driven by theEarth andplanetary science communities, although applications to biological and industrial systems are emerging.[24]
In phases of the meteoritesSemarkona andChervony Kut, a correlation between the concentration of60Ni, thegranddaughter of60Fe, and the abundance of the stable iron isotopes provided evidence for the existence of60Fe at the time offormation of the Solar System. Possibly the energy released by the decay of60Fe, along with that released by26Al, contributed to the remelting anddifferentiation ofasteroids after their formation 4.6 billion years ago. The abundance of60Ni present inextraterrestrial material may bring further insight into the origin and early history of theSolar System.[25]
The most abundant iron isotope56Fe is of particular interest to nuclear scientists because it represents the most common endpoint ofnucleosynthesis.[26] Since56Ni (14alpha particles) is easily produced from lighter nuclei in thealpha process innuclear reactions in supernovae (seesilicon burning process), it is the endpoint of fusion chains insideextremely massive stars. Although adding more alpha particles is possible, but nonetheless the sequence does effectively end at56Ni because conditions in stellar interiors cause the competition betweenphotodisintegration and the alpha process to favor photodisintegration around56Ni.[27][28] This56Ni, 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 thesupernova remnant gas cloud, first to radioactive56Co, and then to stable56Fe. As such, iron is the most abundant element in the core ofred giants, and is the most abundant metal iniron meteorites and in the dense metalcores of planets such asEarth.[29] It is also very common in the universe, relative to other stablemetals of approximately the sameatomic weight.[29][30] Iron is the sixth mostabundant element in theuniverse, and the most commonrefractory element.[31]
Although a further tiny energy gain could be extracted by synthesizing62Ni, which has a marginally higher binding energy than56Fe, 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 than62Ni due to its higher fraction of lighter protons.[32] Hence, elements heavier than iron require asupernova for their formation, involvingrapid neutron capture by starting56Fe nuclei.[29]
In thefar future of the universe, assuming thatproton decay does not occur, coldfusion occurring viaquantum tunnelling would cause the light nuclei in ordinary matter to fuse into56Fe nuclei. Fission andalpha-particle emission would then make heavy nuclei decay into iron, converting all stellar-mass objects to cold spheres of pure iron.[33]
Origin and occurrence in nature
Cosmogenesis
Iron's abundance inrocky planets like Earth is due to its abundant production during the runaway fusion and explosion of typeIa supernovae, which scatters the iron into space.[34][35]
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 ornative iron is rarely found on the surface of the Earth because it tends to oxidize. However, both the Earth'sinner andouter core, which together account for 35% of the mass of the whole Earth, are believed to consist largely of an iron alloy, possibly withnickel. Electric currents in the liquid outer core are believed to be the origin of theEarth's magnetic field. The otherterrestrial planets (Mercury,Venus, andMars) as well as theMoon are believed to have a metallic core consisting mostly of iron. TheM-type asteroids are also believed to be partly or mostly made of metallic iron alloy.
The rareiron meteorites are the main form of natural metallic iron on the Earth's surface. Items made ofcold-worked meteoritic iron have been found in various archaeological sites dating from a time when iron smelting had not yet been developed; and theInuit inGreenland have been reported to use iron from theCape York meteorite for tools and hunting weapons.[36] About 1 in 20meteorites consist of the unique iron-nickel mineralstaenite (35–80% iron) andkamacite (90–95% iron).[37] 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 oxygenfugacity sufficiently for iron to crystallize. This is known astelluric iron and is described from a few localities, such asDisko Island in West Greenland,Yakutia inRussia andBühl inGermany.[38]
Mantle minerals
Ferropericlase(Mg,Fe)O, a solid solution ofpericlase (MgO) andwüstite (FeO), makes up about 20% of the volume of thelower mantle of the Earth, which makes it the second most abundant mineral phase in that region aftersilicate perovskite(Mg,Fe)SiO3; it also is the major host for iron in the lower mantle.[39] At the bottom of thetransition zone of the mantle, the reaction γ-(Mg,Fe)2[SiO4] ↔ (Mg,Fe)[SiO3] + (Mg,Fe)O transformsγ-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.[40]Silicate perovskite may form up to 93% of the lower mantle,[41] and the magnesium iron form,(Mg,Fe)SiO3, is considered to be the most abundantmineral in the Earth, making up 38% of its volume.[42]
While iron is the most abundant element on Earth, most of this iron is concentrated in theinner andouter cores.[43][44] The fraction of iron that is inEarth'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 (afteroxygen,silicon, andaluminium).[45]
Most of the iron in the crust is combined with various other elements to form manyiron minerals. An important class is theiron oxide minerals such ashematite (Fe2O3),magnetite (Fe3O4), andsiderite (FeCO3), which are the majorores of iron. Manyigneous rocks also contain the sulfide mineralspyrrhotite andpentlandite.[46][47] Duringweathering, 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 asiron(III) oxide.[48]
Banded iron formation in McKinley Park, Minnesota
Large deposits of iron arebanded iron formations, a type of rock consisting of repeated thin layers of iron oxides alternating with bands of iron-poorshale andchert. The banded iron formations were laid down in the time between3,700 million years ago and1,800 million years ago.[49][50]
Materials containing finely ground iron(III) oxides or oxide-hydroxides, such asochre, have been used as yellow, red, and brownpigments since pre-historical times. They contribute as well to the color of various rocks andclays, including entire geological formations like thePainted Hills inOregon and theBuntsandstein ("colored sandstone", BritishBunter).[51] ThroughEisensandstein (ajurassic 'iron sandstone', e.g. fromDonzdorf in Germany)[52] andBath stone in the UK, iron compounds are responsible for the yellowish color of many historical buildings and sculptures.[53] The proverbialred color of the surface of Mars is derived from an iron oxide-richregolith.[54]
Significant amounts of iron occur in the iron sulfide mineralpyrite (FeS2), but it is difficult to extract iron from it and it is therefore not exploited.[55] In fact, iron is so common that production generally focuses only on ores with very high quantities of it.[56]
According to theInternational Resource Panel'sMetal 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).[57]
Oceans
Ocean science demonstrated the role of the iron in the ancient seas in both marine biota and climate.[58]
Iron shows the characteristic chemical properties of thetransition metals, namely the ability to form variable oxidation states differing by steps of one and a very large coordination andorganometallic chemistry: indeed, it was the discovery of an iron compound,ferrocene, that revolutionalized the latter field in the 1950s.[59] Iron is sometimes considered as a prototype for the entire block of transition metals, due to its abundance and the immense role it has played in the technological progress of humanity.[60] Its 26 electrons are arranged in theconfiguration [Ar]3d64s2, of which the 3d and 4s electrons are relatively close in energy, and thus a number of electrons can be ionized.[17]
Iron forms compounds mainly in theoxidation states +2 (iron(II), "ferrous") and +3 (iron(III), "ferric"). Iron also occurs inhigher oxidation states, e.g., the purplepotassium ferrate (K2FeO4), which contains iron in its +6 oxidation state. The anion [FeO4]– with iron in its +7 oxidation state, along with an iron(V)-peroxo isomer, has been detected by infrared spectroscopy at 4 K after cocondensation of laser-ablated Fe atoms with a mixture of O2/Ar.[61] Iron(IV) is a common intermediate in many biochemical oxidation reactions.[62][63] Numerousorganoiron 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 ofMössbauer spectroscopy.[64] Manymixed valence compounds contain both iron(II) and iron(III) centers, such asmagnetite andPrussian blue (Fe4(Fe[CN]6)3).[63] The latter is used as the traditional "blue" inblueprints.[65]
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.[11] Ruthenium exhibits an aqueous cationic chemistry in its low oxidation states similar to that of iron, but osmium does not, favoring high oxidation states in which it forms anionic complexes.[11] In the second half of the 3d transition series, vertical similarities down the groups compete with the horizontal similarities of iron with its neighborscobalt andnickel in the periodic table, which are also ferromagnetic atroom temperature and share similar chemistry. As such, iron, cobalt, and nickel are sometimes grouped together as theiron triad.[60]
Unlike many other metals, iron does not form amalgams withmercury. As a result, mercury is traded in standardized 76 pound flasks (34 kg) made of iron.[66]
Iron is by far the most reactive element in its group; it ispyrophoric when finely divided and dissolves easily in dilute acids, giving Fe2+. However, it does not react with concentratednitric acid and other oxidizing acids due to the formation of an impervious oxide layer, which can nevertheless react withhydrochloric acid.[11] High-purity iron, calledelectrolytic iron, is considered to be resistant to rust, due to its oxide layer.
Binary compounds
Oxides and sulfides
Ferrous or iron(II) oxide,FeO
Ferric or iron(III) oxideFe2O3
Ferrosoferric or iron(II,III) oxideFe3O4
Iron forms variousoxide and hydroxide compounds; the most common areiron(II,III) oxide (Fe3O4), andiron(III) oxide (Fe2O3).Iron(II) oxide also exists, though it is unstable at room temperature. Despite their names, they are actually allnon-stoichiometric compounds whose compositions may vary.[67] These oxides are the principal ores for the production of iron (seebloomery and blast furnace). They are also used in the production offerrites, usefulmagnetic storage media in computers, and pigments. The best known sulfide isiron pyrite (FeS2), also known as fool's gold owing to its golden luster.[63] It is not an iron(IV) compound, but is actually an iron(II)polysulfide containing Fe2+ andS2− 2 ions in a distortedsodium chloride structure.[67]
The binary ferrous and ferrichalides are well-known. The ferrous halides typically arise from treating iron metal with the correspondinghydrohalic acid to give the corresponding hydrated salts.[63]
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.[68]
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−:[68]
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 ofiron pentacarbonyl withiodine andcarbon monoxide in the presence ofhexane and light at the temperature of −20 °C, with oxygen and water excluded.[68] Complexes of ferric iodide with some soft bases are known to be stable compounds.[69][70]
Solution chemistry
Comparison of colors of solutions of ferrate (left) andpermanganate (right)
As pH rises above 0 the above yellow hydrolyzed species form and as it rises above 2–3, reddish-brown hydrousiron(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-metalcharge 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.[71] 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 whencarbonate anions are added, which instead results in whiteiron(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 formiron(III) oxide that accounts for the brown deposits present in a sizeable number of streams.[72]
Coordination compounds
Due to its electronic structure, iron has a very large coordination and organometallic chemistry.
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 ofchromium(III) with the exception of iron(III)'s preference forO-donor instead ofN-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 forphenols orenols. For example, in theferric chloride test, used to determine the presence of phenols,iron(III) chloride reacts with a phenol to form a deep violet complex:[71]
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 thespectrochemical series such ascyanide. An example of a low-spin iron(III) complex is [Fe(CN)6]3−. Iron shows a great variety of electronicspin states, including every possible spin quantum number value for a d-block element from 0 (diamagnetic) to5⁄2 (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.[67]
Iron(II) complexes are less stable than iron(III) complexes but the preference forO-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.[72]
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) withpotassium ferricyanide andpotassium ferrocyanide to form Prussian blue.[63]
Another old example of an organoiron compound isiron pentacarbonyl, Fe(CO)5, in which a neutral iron atom is bound to the carbon atoms of fivecarbon monoxide molecules. The compound can be used to makecarbonyl iron powder, a highly reactive form of metallic iron.Thermolysis of iron pentacarbonyl givestriiron 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.[78]
Structural formula of ferrocene and a powdered sample
A landmark in this field was the discovery in 1951 of the remarkably stablesandwich compoundferroceneFe(C5H5)2, by Pauson and Kealy[79] and independently by Miller and colleagues,[80] whose surprising molecular structure was determined only a year later byWoodward andWilkinson[81] andFischer.[82]Ferrocene is still one of the most important tools and models in this class.[83]
The iron compounds produced on the largest scale in industry areiron(II) sulfate (FeSO4·7H2O) andiron(III) chloride (FeCl3). The former is one of the most readily available sources of iron(II), but is less stable to aerial oxidation thanMohr's salt ((NH4)2Fe(SO4)2·6H2O). Iron(II) compounds tend to be oxidized to iron(III) compounds in the air.[63]
Iron is one of the elements undoubtedly known to the ancient world.[85] It has been worked, orwrought, 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.[86] 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 fromGreenland. The iron edge covers anarwhal tusk harpoon using meteorite iron from theCape York meteorite, one of the largest iron meteorites known.
Beads made frommeteoric iron in 3500 BC or earlier were found inGerzeh, Egypt byG. A. Wainwright.[87] 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.[87] For example, adagger made of meteoric iron was found in the tomb ofTutankhamun, containing similar proportions of iron, cobalt, and nickel to a meteorite discovered in the area, deposited by an ancient meteor shower.[88][89][90] Items that were likely made of iron by Egyptians date from 3000 to 2500 BC.[86]
Meteoritic iron is comparably soft and ductile and easilycold forged but may get brittle when heated because of thenickel content.[91]
The symbol forMars has been used since antiquity to represent iron.Theiron pillar of Delhi is an example of the iron extraction and processing methodologies of early India.
The first iron production started in theMiddle Bronze Age, but it took several centuries before iron displaced bronze. Samples ofsmelted iron fromAsmar, Mesopotamia and Tall Chagar Bazaar in northern Syria were made sometime between 3000 and 2700 BC.[92] TheHittites established an empire in north-centralAnatolia 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.[93] TheHittites 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.[92] The subsequent period is called theIron Age.
Artifacts of smelted iron are found inIndia dating from 1800 to 1200 BC,[94] and in theLevant from about 1500 BC (suggesting smelting inAnatolia or theCaucasus).[95][96] Alleged references (comparehistory of metallurgy in South Asia) to iron in the IndianVedas have been used for claims of a very early usage of iron in India respectively to date the texts as such. Therigveda termayas (metal) refers to copper, while iron which is called asśyāma ayas, literally "black copper", first is mentioned in the post-rigvedicAtharvaveda.[97]
Some archaeological evidence suggests iron was smelted inZimbabwe and southeast Africa as early as the eighth century BC.[98] Iron working was introduced toGreece in the late 11th century BC, from which it spread quickly throughout Europe.[99]
Iron sickle from Ancient Greece
The spread of ironworking in Central and Western Europe is associated withCeltic expansion. According toPliny the Elder, iron use was common in theRoman era.[87] In the lands of what is now consideredChina, iron appears approximately 700–500 BC.[100] Iron smelting may have been introduced into China through Central Asia.[101] The earliest evidence of the use of ablast furnace in China dates to the 1st century AD,[102] and cupola furnaces were used as early as theWarring States period (403–221 BC).[103] Usage of the blast and cupola furnace remained widespread during theTang andSong dynasties.[104]
During the Industrial Revolution in Britain,Henry Cort began refining iron frompig iron towrought iron (or bar iron) using innovative production systems. In 1783 he patented thepuddling process for refining iron ore. It was later improved by others, includingJoseph Hall.[105]
Cast iron was first produced in China during 5th century BC,[106] but was hardly in Europe until the medieval period.[107][108] The earliest cast iron artifacts were discovered by archaeologists in what is now modernLuhe County,Jiangsu in China. Cast iron was used inancient China for warfare, agriculture, and architecture.[109] During themedieval period, means were found in Europe of producing wrought iron from cast iron (in this context known aspig iron) usingfinery forges. For all these processes,charcoal was required as fuel.[110]
Medievalblast furnaces were about 10 feet (3.0 m) tall and made of fireproof brick; forced air was usually provided by hand-operated bellows.[108] 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.[110]
In 1709,Abraham Darby I established acoke-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 theIndustrial 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.[92]
Since iron was becoming cheaper and more plentiful, it also became a major structural material following the building of the innovativefirst 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 insteam engines.[110] Railways have been central to the formation of modernity and ideas of progress[111] and various languages refer to railways asiron road (e.g. Frenchchemin de fer, GermanEisenbahn, Turkishdemiryolu, Russianжелезная дорога,Chinese, Japanese, and Korean 鐵道, Vietnameseđường sắt).
Steel (with smaller carbon content than pig iron but more than wrought iron) was first produced in antiquity by using abloomery. Blacksmiths inLuristan in western Persia were making good steel by 1000 BC.[92] Then improved versions,Wootz steel by India andDamascus steel were developed around 300 BC and AD 500 respectively. These methods were specialized, and so steel did not become a major commodity until the 1850s.[112]
New methods of producing it bycarburizing bars of iron in thecementation process were devised in the 17th century. In theIndustrial 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 in large quantities.[113]
Foundations of modern chemistry
In 1774,Antoine Lavoisier used the reaction of water steam with metallic iron inside an incandescent iron tube to producehydrogen in his experiments leading to the demonstration of theconservation of mass, which was instrumental in changing chemistry from a qualitative science to a quantitative one.[114]
Symbolic role
"Ich gab Gold für Eisen" – "I gave gold for iron".German-American brooch from WWI.
Iron plays a certain role in mythology and has found various usageas a metaphor and infolklore. TheGreek poetHesiod'sWorks and Days (lines 109–201) lists differentages of man named after metals like gold, silver, bronze and iron to account for successive ages of humanity.[115] The Iron Age was closely related with Rome, and in Ovid'sMetamorphoses
The Virtues, in despair, quit the earth; and the depravity of man becomes universal and complete. Hard steel succeeded then.
An example of the importance of iron's symbolic role may be found in theGerman Campaign of 1813.Frederick William III commissioned then the firstIron Cross as military decoration.Berlin iron jewellery reached its peak production between 1813 and 1815, when the Prussianroyal family urged citizens to donate gold and silver jewellery for military funding. The inscriptionIch gab Gold für Eisen (I gave gold for iron) was used as well in later war efforts.[116]
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.[48] Another method is electrolysis of ferrous chloride onto an iron cathode.[117]
Nowadays, the industrial production of iron or steel consists of two main stages. In the first stage, iron ore isreduced withcoke in ablast furnace, and the molten metal is separated from gross impurities such assilicate 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.[119] Other metals can be added at this stage to formalloy steels.
The blast furnace is loaded with iron ores, usuallyhematiteFe2O3 ormagnetiteFe3O4, along with coke (coal that has been separately baked to remove volatile components) andflux (limestone ordolomite). "Blasts" of air pre-heated to 900 °C (sometimes with oxygen enrichment) is blown through the mixture, in sufficient amount to turn the carbon intocarbon monoxide:[119]
2 C + O2 → 2 CO
This reaction raises the temperature to about 2000 °C. The carbon monoxide reduces the iron ore to metallic iron:[119]
Fe2O3 + 3 CO → 2 Fe + 3 CO2
Some iron in the high-temperature lower region of the furnace reacts directly with the coke:[119]
2 Fe2O3 + 3 C → 4 Fe + 3 CO2
The flux removes silicaceous minerals in the ore, which would otherwise clog the furnace: The heat of the furnace decomposes the carbonates tocalcium oxide, which reacts with any excesssilica to form aslag composed ofcalcium silicateCaSiO3 or other products. At the furnace's temperature, the metal and the slag are both molten. They collect at the bottom as two immiscible liquid layers (with the slag on top), that are then easily separated.[119] The slag can be used as a material inroad construction or to improve mineral-poor soils foragriculture.[108]
Steelmaking thus remains one of the largest industrial contributors of CO2 emissions in the world.[120]
17th century Chinese illustration of workers at a blast furnace, making wrought iron from pig iron[121]
The pig iron produced by the blast furnace process contains up to 4–5% carbon (by mass), with small amounts of other impurities like sulfur, magnesium, phosphorus, and manganese. This high level of carbon makes it relatively weak and brittle. Reducing the amount of carbon to 0.002–2.1% producessteel, which may be up to 1000 times harder than pure iron. A great variety of steel articles can then be made bycold working,hot rolling,forging,machining, etc. Removing the impurities from pig iron, but leaving 2–4% carbon, results incast iron, which is cast byfoundries into articles such as stoves, pipes, radiators, lamp-posts, and rails.[119]
Steel products often undergo variousheat treatments after they are forged to shape.Annealing consists of heating them to 700–800 °C for several hours and then gradual cooling. It makes the steel softer and more workable.[122]
This heap of iron ore pellets will be used in steel production.
A pot of molten iron being used to make steel
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.[108] Two main reactions comprise the direct reduction process:
Natural gas is partially oxidized (with heat and a catalyst):[108]
2 CH4 + O2 → 2 CO + 4 H2
Iron ore is then treated with these gases in a furnace, producing solid sponge iron:[108]
Ignition of a mixture of aluminium powder and iron oxide yields metallic iron via thethermite reaction:
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, includingfinery forges,puddling furnaces,Bessemer converters,open hearth furnaces,basic oxygen furnaces, andelectric 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.[110]
Molten oxide electrolysis
Molten oxide electrolysis (MOE) useselectrolysis of molten iron oxide to yield metallic iron. It is studied in laboratory-scale experiments and is proposed as a method for industrial iron production that has no direct emissions of carbon dioxide. It uses a liquid iron cathode, an anode formed from an alloy of chromium, aluminium and iron,[123] and 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. Oxygen gas is produced in addition to liquid iron. The only carbon dioxide emissions come from anyfossil fuel-generated electricity used to heat and reduce the metal.[124][125][126]
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 andmachine 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.[129]
Mechanical properties
The mechanical properties of iron and its alloys are extremely relevant to their structural applications. Those properties can be evaluated in various ways, including theBrinell test, theRockwell test and theVickers hardness test.
The properties of pure iron are often used to calibrate measurements or to compare tests.[128][130] However, the mechanical properties of iron are significantly affected by the sample's purity: pure, single crystals of iron are actually softer than aluminium,[127] and the purest industrially produced iron (99.99%) has a hardness of 20–30 Brinell.[131] The pure iron (99.9%～99.999%), especially calledelectrolytic iron, is industrially produced byelectrolytic refining.
An increase in the carbon content will cause a significant increase in the hardness and tensile strength of iron. Maximum hardness of65 Rc is achieved with a 0.6% carbon content, although the alloy has low tensile strength.[132] Because of the softness of iron, it is much easier to work with than its heaviercongenersruthenium andosmium.[17]
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).[133]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 ofstainless steel used for making cutlery, and hospital and food-service equipment.[21]
Commercially available iron is classified based on purity and the abundance of additives.Pig iron has 3.5–4.5% carbon[134] and contains varying amounts of contaminants such assulfur, silicon andphosphorus. 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 ofmanganese.[119] Pig iron has amelting 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.[11] Its mechanical properties vary greatly and depend on the form the carbon takes in the alloy.[17]
"White" cast irons contain their carbon in the form ofcementite, or iron carbide (Fe3C).[17] 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 calledpearlite for its appearance. Rapid cooling, on the other hand, does not allow time for this separation and creates hard and brittlemartensite. 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.[17]
Ingray iron the carbon exists as separate, fine flakes ofgraphite, and also renders the material brittle due to the sharp edged flakes of graphite that producestress concentration sites within the material.[135] A newer variant of gray iron, referred to asductile iron, is specially treated with trace amounts ofmagnesium to alter the shape of graphite to spheroids, or nodules, reducing the stress concentrations and vastly increasing the toughness and strength of the material.[135]
Wrought iron contains less than 0.25% carbon but large amounts ofslag that give it a fibrous characteristic.[134] Wrought iron is more corrosion resistant than steel. It has been almost completely replaced bymild steel, which corrodes more readily than wrought iron, but is cheaper and more widely available.Carbon steel contains 2.0% carbon or less,[136] with small amounts ofmanganese,sulfur,phosphorus, and silicon.Alloy steels contain varying amounts of carbon as well as other metals, such aschromium,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, isstainless 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.[136][137][138]
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.[139]
The main disadvantage of iron and steel is that pure iron, and most of its alloys, suffer badly fromrust if not protected in some way, a cost amounting to over 1% of the world's economy.[140]Painting,galvanization,passivation, plastic coating andbluing are all used to protect iron from rust by excludingwater and oxygen or bycathodic protection. The mechanism of the rusting of iron is as follows:[140]
The electrolyte is usuallyiron(II) sulfate in urban areas (formed when atmosphericsulfur dioxide attacks iron), and salt particles in the atmosphere in seaside areas.[140]
Catalysts and reagents
Because Fe is inexpensive and nontoxic, much effort has been devoted to the development of Fe-based catalysts andreagents. Iron is however less common as a catalyst in commercial processes than more expensive metals.[141] In biology, Fe-containing enzymes are pervasive.[142]
Iron(III) chloride finds use in water purification andsewage treatment, in the dyeing of cloth, as a coloring agent in paints, as an additive in animal feed, and as anetchant forcopper in the manufacture ofprinted circuit boards.[145] It can also be dissolved in alcohol to form tincture of iron, which is used as a medicine to stop bleeding incanaries.[146]
Iron(II) sulfate is used as a precursor to other iron compounds. It is also used toreduce chromate in cement. It is used to fortify foods and treatiron 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.[145]
Simplified structure ofHeme B; in the protein additionalligand(s) are attached to Fe.
Examples of iron-containing proteins in higher organisms include hemoglobin,cytochrome (seehigh-valent iron), andcatalase.[10][151] The average adult human contains about 0.005% body weight of iron, or about four grams, of which three quarters is in hemoglobin—a level that remains constant despite only about one milligram of iron being absorbed each day,[150] because the human body recycles its hemoglobin for the iron content.[152]
Microbial growth may be assisted by oxidation of iron(II) or by reduction of iron(III).[153]
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.[10] In particular, bacteria have evolved very high-affinitysequestering agents calledsiderophores.[154][155][156]
After uptake in humancells, iron storage is precisely regulated.[10][157] A major component of this regulation is the proteintransferrin, which binds iron ions absorbed from theduodenum and carries it in theblood to cells.[10][158] Transferrin contains Fe3+ in the middle of a distorted octahedron, bonded to one nitrogen, three oxygens and a chelatingcarbonate anion that traps the Fe3+ ion: it has such a highstability 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+ to Fe2+ and stored asferritin to be incorporated into hemoglobin.[150]
Hemoglobin is an oxygen carrier that occurs inred blood cells and contributes their color, transporting oxygen in the arteries from the lungs to the muscles where it is transferred tomyoglobin, which stores it until it is needed for the metabolic oxidation ofglucose, generating energy.[10] Here the hemoglobin binds tocarbon dioxide, produced when glucose is oxidized, which is transported through the veins by hemoglobin (predominantly asbicarbonate anions) back to the lungs where it is exhaled.[150] In hemoglobin, the iron is in one of fourheme groups and has six possible coordination sites; four are occupied by nitrogen atoms in aporphyrin ring, the fifth by animidazole nitrogen in ahistidine 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.[150] When hemoglobin is not attached to oxygen (and is then called deoxyhemoglobin), the Fe2+ ion at the center of theheme group (in the hydrophobic protein interior) is in ahigh-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.[150]
When deoxyhemoglobin picks up an oxygen molecule, this histidine residue moves away and returns once the oxygen is securely attached to form ahydrogen 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.[150] 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.[150] 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.[150] As described by theBohr effect (named afterChristian Bohr, the father ofNiels Bohr), the oxygen affinity of hemoglobin diminishes in the presence of carbon dioxide.[150]
Carbon monoxide andphosphorus 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 ascarboxyhemoglobin. This effect also plays a minor role in the toxicity ofcyanide, but there the major effect is by far its interference with the proper functioning of the electron transport proteincytochrome a.[150] 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 amethionine 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.[150] 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 inadenosine 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:[150]
Although the heme proteins are the most important class of iron-containing proteins, theiron–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 isrubredoxin, which has only one iron atom coordinated to four sulfur atoms fromcysteine residues in the surrounding peptide chains. Another important class of iron–sulfur proteins is theferredoxins, which have multiple iron atoms. Transferrin does not belong to either of these classes.[150]
The ability of seamussels to maintain their grip on rocks in the ocean is facilitated by their use oforganometallic iron-based bonds in their protein-richcuticles. Based on synthetic replicas, the presence of iron in these structures increasedelastic modulus 770 times,tensile strength 58 times, andtoughness 92 times. The amount of stress required to permanently damage them increased 76 times.[163]
Iron provided bydietary supplements is often found asiron(II) fumarate, althoughiron(II) sulfate is cheaper and is absorbed equally well.[145] Elemental iron, or reduced iron, despite being absorbed at only one-third to two-thirds the efficiency (relative to iron sulfate),[165] is often added to foods such as breakfast cereals or enriched wheat flour. Iron is most available to the body whenchelated to amino acids[166] and is also available for use as a commoniron supplement.Glycine, the least expensive amino acid, is most often used to produce iron glycinate supplements.[167]
Dietary recommendations
The U.S. Institute of Medicine (IOM) updated Estimated Average Requirements (EARs) and Recommended Dietary Allowances (RDAs) for iron in 2001.[10] The current EAR for iron for women ages 14–18 is 7.9 mg/day, 8.1 mg/day for ages 19–50 and 5.0 mg/day thereafter (postmenopause). For men, the EAR is 6.0 mg/day for ages 19 and up. The RDA is 15.0 mg/day for women ages 15–18, 18.0 mg/day for ages 19–50 and 8.0 mg/day thereafter. For men, 8.0 mg/day for ages 19 and up. RDAs are higher than EARs so as to identify amounts that will cover people with higher-than-average requirements. RDA for pregnancy is 27 mg/day and, for lactation, 9 mg/day.[10] For children ages 1–3 years 7 mg/day, 10 mg/day for ages 4–8 and 8 mg/day for ages 9–13. As for safety, the IOM also setsTolerable upper intake levels (ULs) for vitamins and minerals when evidence is sufficient. In the case of iron, the UL is set at 45 mg/day. Collectively the EARs, RDAs and ULs are referred to asDietary Reference Intakes.[168]
TheEuropean Food Safety Authority (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR. AI and UL are defined the same as in theUnited States. For women the PRI is 13 mg/day ages 15–17 years, 16 mg/day for women ages 18 and up who are premenopausal and 11 mg/day postmenopausal. For pregnancy and lactation, 16 mg/day. For men the PRI is 11 mg/day ages 15 and older. For children ages 1 to 14, the PRI increases from 7 to 11 mg/day. The PRIs are higher than the U.S. RDAs, with the exception of pregnancy.[169] The EFSA reviewed the same safety question did not establish a UL.[170]
Infants may require iron supplements if they are bottle-fed cow's milk.[171] Frequentblood donors are at risk of low iron levels and are often advised to supplement their iron intake.[172]
For U.S. food and dietary supplement labeling purposes, the amount in a serving is expressed as a percent of Daily Value (%DV). For iron labeling purposes, 100% of the Daily Value was 18 mg, and as of May 27, 2016[update] remained unchanged at 18 mg.[173][174] A table of the old and new adult daily values is provided atReference Daily Intake.
Iron deficiency is the most commonnutritional deficiency in the world.[10][175][176][177] When loss of iron is not adequately compensated by adequate dietary iron intake, a state oflatent iron deficiency occurs, which over time leads toiron-deficiency anemia if left untreated, which is characterised by an insufficient number of red blood cells and an insufficient amount of hemoglobin.[178] 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.[179]
The brain is resistant to acute iron deficiency due to the slow transport of iron through the blood brain barrier.[180] 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.[181][182] In the brain, iron plays a role in oxygen transport, myelin synthesis, mitochondrial respiration, and as a cofactor for neurotransmitter synthesis and metabolism.[183] Animal models of nutritional iron deficiency report biomolecular changes resembling those seen in Parkinson's and Huntington's disease.[184][185] However, age-related accumulation of iron in the brain has also been linked to the development of Parkinson's.[186]
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.[187] Regulation of iron uptake is impaired in some people as a result of agenetic defect that maps to the HLA-H gene region onchromosome 6 and leads to abnormally low levels ofhepcidin, a key regulator of the entry of iron into the circulatory system in mammals.[188] In these people, excessive iron intake can result iniron overload disorders, known medically ashemochromatosis.[10] 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 fromiron deficiency and have consulted a doctor. Hemochromatosis is estimated to be the cause of 0.3–0.8% of all metabolic diseases of Caucasians.[189]
Overdoses of ingested iron can cause excessive levels of free iron in the blood. High blood levels of free ferrous iron react withperoxides to produce highly reactivefree radicals that can damageDNA,proteins,lipids, and other cellular components. Iron toxicity occurs when the cell contains free iron, which generally occurs when iron levels exceed the availability oftransferrin to bind the iron. Damage to the cells of thegastrointestinal tract can also prevent them from regulating iron absorption, leading to further increases in blood levels. Iron typically damages cells in theheart,liver and elsewhere, causing adverse effects that includecoma,metabolic acidosis,shock,liver failure,coagulopathy, long-term organ damage, and even death.[190] Humans experience iron toxicity when the iron exceeds 20 milligrams for every kilogram of body mass; 60 milligrams per kilogram is considered alethal dose.[191] Overconsumption of iron, often the result of children eating large quantities offerrous sulfate tablets intended for adult consumption, is one of the most common toxicological causes of death in children under six.[191] TheDietary 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.[192]
The medical management of iron toxicity is complicated, and can include use of a specificchelating agent calleddeferoxamine to bind and expel excess iron from the body.[190][193][194]
ADHD
Some research has suggested that lowthalamic iron levels may play a role in the pathophysiology ofADHD.[195] Some researchers have found that iron supplementation can be effective especially in theinattentive subtype of the disorder.[196]
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.[197]
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.[198] People havingchemotherapy may develop iron deficiency andanemia, for whichintravenous iron therapy is used to restore iron levels.[199] Iron overload, which may occur from high consumption of red meat,[10] may initiatetumor growth and increase susceptibility to cancer onset,[199] particularly forcolorectal cancer.[10]
Marine systems
Iron plays an essential role in marine systems and can act as a limiting nutrient for planktonic activity.[200] Because of this, too much of a decrease in iron may lead to a decrease in growth rates in phytoplanktonic organisms such as diatoms.[201] Iron can also be oxidized by marine microbes under conditions that are high in iron and low in oxygen.[202]
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.[203] 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.[204] The iron cycle can fluctuate the forms of iron from aqueous to particle forms altering the availability of iron to primary producers.[205] Increased light and warmth increases the amount of iron that is in forms that are usable by primary producers.[206]
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