Iron

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"Fe" redirects here. For other uses, see Fe (disambiguation).
This article is about the chemical element. For other uses, see Iron (disambiguation).
manganese ← iron → cobalt
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Fe
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Ru
Appearance
lustrous metallic with a grayish tinge
250px
250px
Spectral lines of Iron
General properties
Name, symbol, number iron, Fe, 26
Pronunciation US /aɪ.ərn/; UK /ˈaɪərn/
Element category transition metal
Group, period, block 8, 4, d
Standard atomic weight 55.845(2) g·mol−1
Electron configuration [Ar] 3d6 4s2
Electrons per shell 2, 8, 14, 2 (Image)
Physical properties
Phase solid
Density (near r.t.) 7.874 g·cm−3
Liquid density at m.p. 6.98 g·cm−3
Melting point 1811 K, 1538 °C, 2800 °F
Boiling point 3134 K, 2862 °C, 5182 °F
Heat of fusion 13.81 kJ·mol−1
Heat of vaporization 340 kJ·mol−1
Specific heat capacity (25 °C) 25.10 J·mol−1·K−1
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 1728 1890 2091 2346 2679 3132
Atomic properties
Electronegativity 1.83 (Pauling scale)
Ionization energies
(more) 		1st: 762.5 kJ·mol−1
2nd: 1561.9 kJ·mol−1
3rd: 2957 kJ·mol−1
Atomic radius 126 pm
Covalent radius 132±3 (low spin), 152±6 (high spin) pm
Miscellanea
Crystal structure body-centered cubic
Magnetic ordering ferromagnetic
1043 K
Electrical resistivity (20 °C) 96.1 nΩ·m
Thermal conductivity (300 K) 80.4 W·m−1·K−1
Thermal expansion (25 °C) 11.8 µm·m−1·K−1
Speed of sound (thin rod) (r.t.) (electrolytic)
5120 m·s−1
Young's modulus 211 GPa
Shear modulus 82 GPa
Bulk modulus 170 GPa
Poisson ratio 0.29
Mohs hardness 4
Vickers hardness 608 MPa
Brinell hardness 490 MPa
CAS registry number 7439-89-6
Most stable isotopes
Main article: Isotopes of iron
iso NA half-life DM DE (MeV) DP
54Fe 5.8% >3.1×1022y 2ε capture  ? 54Cr
55Fe syn 2.73 y ε capture 0.231 55Mn
56Fe 91.72% 56Fe is stable with 30 neutrons
57Fe 2.2% 57Fe is stable with 31 neutrons
58Fe 0.28% 58Fe is stable with 32 neutrons
59Fe syn 44.503 d β− 1.565 59Co
60Fe syn 2.6×106 y β− 3.978 60Co

Iron (11px /ˈaɪ.ərn/ or /ˈaɪrn/) is a chemical element with the symbol Fe (Latin: ferrum) and atomic number 26. It is a metal in the first transition series. Like other Group 8 elements, it exists in a wide range of oxidation states. Iron and iron alloys (steels) are by far the most common metals and the most common ferromagnetic materials in everyday use. Fresh iron surfaces appear lustrous silvery-gray, but oxidize in air.

Iron is the most common element in the Earth, and the fourth most common in the Earth's crust. It is produced as a result of fusion in high-mass stars, and it is the heaviest stable element produced by stellar fusion because the fusion of iron is the last nuclear fusion reaction that is exothermic. Iron is the most widely used metal, and iron compounds, which include ferrous and ferric compounds, have several uses as well.

Iron has been used since ancient times, though not as early as bronze or the other copper related alloys. Iron is ubiquitous in modern life; it is used primarily for its structural strength. Pure iron is soft (softer than aluminium), but the material is significantly strengthened by addition of minute amounts of impurities, such as carbon. Alloying iron with appropriate small amounts (up to a few per cent) of other metals and carbon produces steel, which can be 1,000 times harder than pure iron. Iron is smelted in a blast furnace, where ore is reduced by coke to metallic iron.

Elemental iron is reactive; it oxidizes in air to give iron oxides, also known as rust. The rusting of iron and iron alloys is undesirable, and has a major economic impact. Unlike many other metals which form passivating oxide layers, iron oxides occupy more volume than iron itself. Thus, iron oxides flake off and expose fresh surfaces for corrosion. Iron oxide mixed with aluminium powder can be ignited to create a thermite reaction, used in welding and purifying ores.

Iron exists from oxidation state −2 to + 6, although +2 and +3 are the most common. It forms binary compounds with the halogens and the chalcogens. Among its organometallic compounds, ferrocene was the first sandwich compound discovered. Iron plays an important role in biology, forming complexes with dioxygen as hemoglobin and myoglobin; these two compounds are common oxygen transport proteins in vertebrates.

Characteristics

Mechanical properties

Characteristic values of tensile strength (TS) and Brinell hardness (BH) of different forms of iron.[1][2]
Material TS (MPa) BH (Brinell)
Iron whiskers 11000
Ausformed (hardened) steel 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
Iron containing dissolved carbon 140 40
Single crystal of pure iron 10 3

Mechanical properties of iron and its alloys are evaluated using a variety of tests, such as the Brinell test, Rockwell test, or tensile strength tests, among others; the results on iron are so consistent that iron is often used to calibrate measurements or to relate the results of one test to another.[2][3] Those measurements reveal that mechanical properties of iron crucially depend on purity: Purest research-purpose single crystals of iron are softer than aluminium. Addition of only 10 parts per million of carbon doubles their strength.[1] The hardness increases rapidly with carbon content up to 0.2% and saturates at ~0.6%.[4] The purest industrially produced iron (about 99.99% purity) has a hardness of 20–30 Brinell.[5]

Allotropes

Main article: Allotropes of iron

Iron represents perhaps the best-known example of allotropy in a metal. There are three allotropic forms of iron, known as α, γ and δ.

As molten iron cools down it crystallizes at 1538 °C into its δ allotrope, which has a body-centered cubic (bcc) crystal structure. As it cools further its crystal structure changes to face-centered cubic (fcc) at 1394 °C, when it is known as γ-iron, or austenite. At 912 °C the crystal structure again becomes bcc as α-iron, or ferrite, is formed, and 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 are in the same direction; the neighboring domains point in various directions and thus cancel out. In magnetized iron, the electronic spins of all the domains are aligned, so that the magnetic effects of neighboring domains reinforce each other. Although each domain contains billions of atoms, they are very small, about 10 microns across.[6]

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

Alpha iron, also known as ferrite, is the most stable form of iron at normal temperatures. It is a fairly soft metal that can dissolve only a small concentration of carbon (no more than 0.021% by mass at 910 °C).[7]

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.[6]

Isotopes

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. The nuclide 54Fe is predicted to undergo double beta decay, but this process had never been observed experimentally for these nuclei, and only the lower limit on the half-life was established: T1/2>3.1×1022 years. 60Fe is an extinct radionuclide of long half-life (2.6 million years).[8]

Much of the past work on measuring the isotopic composition of Fe has focused on determining 60Fe variations due to processes accompanying nucleosynthesis (i.e., meteorite studies) and ore formation. In the last decade however, advances in mass spectrometry technology have allowed the detection and quantification of minute, naturally occurring variations in the ratios of the stable isotopes of iron. Much of this work has been driven by the Earth and planetary science communities, although applications to biological and industrial systems are beginning to emerge.[9]

The most abundant iron isotope 56Fe is of particular interest to nuclear scientists as it represents the most stable nuclide possible. It is impossible to perform fission or fusion on 56Fe and still liberate energy. 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 therefore made in quantity in these stars, but soon decays by two successive positron emissions within supernova decay products in the supernova remnant gas cloud, to first radioactive cobalt-56, and then stable iron-56. This last nuclide is therefore common in the universe, relative to other stable metals of approximately the same atomic weight.

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

Nuclei of iron atoms have some of the highest binding energies per nucleon, surpassed only by the nickel isotope 62Ni. This is formed by nuclear fusion in stars. Although a further tiny energy gain could be extracted by synthesizing 62Ni, conditions in stars are unsuitable for this process to be favored. Elemental distribution on Earth greatly favors iron over nickel, and also presumably in supernova element production.[10]

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.

Nucleosynthesis

Iron is created in extremely large, extremely hot (over 2.5 billion kelvin) stars, in a process called the silicon burning process. It is the last element to be produced in this manner. The process starts with the second largest stable nucleus created by silicon burning: 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. The nickel then decays to unstable cobalt, which finally decays to stable iron-56. The iron can no longer be fused with other elements. As the star's core fills with iron, it begins to cool until there is no longer enough energy to maintain its size. It then collapses and the result is a supernova. Supernovas also create additional forms of stable iron via the r-process.

Occurrence

See also Category: Iron minerals

Planetary occurrence

File:Widmanstatten-patterns-3.jpg
Iron meteorites of similar composition of Earth's inner and outer core

Iron is the sixth most abundant element in the Universe, formed as the final step of nucleosynthesis, by silicon fusing in massive stars. Metallic 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.[11][12] Most of the iron in the crust is found combined with oxygen as iron oxide minerals such as hematite and magnetite. 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, either magnetite (Fe3O4) or hematite (Fe2O3), alternating with bands of iron-poor shale and chert. The banded iron formations are common in the time between 3,700 million years ago and 1,800 million years ago[13][14]

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.[15] It was proven by Mössbauer spectroscopy that the red color of the surface of Mars is derived from an iron oxide-rich regolith.[16]

Biological role

Iron is abundant in biology. Iron-proteins are found in all living organisms, ranging from the evolutionarily primitive archaea to humans. The color of blood is due to the hemoglobin, an iron-containing protein. As illustrated by hemoglobin, iron often is bound to cofactors, e.g. in hemes. The iron-sulfur clusters are pervasive and include nitrogenase, the enzymes responsible for biological nitrogen fixation.(see #Biological role). Influential theories of evolution have invoked a role for iron sulfides, iron-sulfur world theory.

Chemistry and compounds

See also Category: Iron compounds

Iron's oxidation states
Oxidation state Representative compound
−2 Disodium tetracarbonylferrate (Collman's reagent)
−1
0 Iron pentacarbonyl
1 Cyclopentadienyliron dicarbonyl dimer ("Fp2")
2 Ferrous sulfate, ferrocene
3 Ferric chloride, ferrocenium tetrafluoroborate
4 Barium ferrate(IV)
5
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. Iron(IV) is a common intermediate in many in biochemical oxidation reactions.[17][18] 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.[19] 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).[18] The latter is used as the traditional "blue" in blueprints.[20]

File:Iron(III) chloride hexahydrate.jpg
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.[18]

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.[21]

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.[18]

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 binary halogen acid to give the corresponding hydrated salts.[18]

Fe + 2 HX → FeX2 + H2

Iron reacts with fluorine, chlorine, and bromine to give the corresponding ferric halides, ferric chloride being the most common:

2 Fe + 3 X2 → 2 FeX3 (X = F, Cl, Br)

Complexation and organometallic compounds

See also: organoiron chemistry
File:Prussian blue.jpg
Prussian blue

Several cyanide complexes are known. The most famous example is Prussian blue, (Fe4(Fe[CN]6)3). Potassium ferricyanide and potassium ferrocyanide are also known; the formation of Prussian blue upon reaction with iron(II) and iron(III) respectively forms the basis of a "wet" chemical test.[18] Prussian blue is also used as an antidote for thallium and radioactive caesium poisoning.[22][23] Prussian blue can be used in laundry bluing to correct the yellowish tint left by ferrous salts in water.[24]

File:Ferrocene-2D.png
Ferrocene

Several carbonyl compounds of iron are known. 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.[25]

Ferrocene is an extremely stable complex. The first sandwich compound, it contains an iron(II) center with two cyclopentadienyl ligands bonded through all ten carbon atoms. This arrangement was a shocking novelty when it was first discovered,[26] but the discovery of ferrocene has led to a new branch of organometallic chemistry. 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.[27]

Bioinorganic compounds

The most famous bioinorganic compounds of iron are heme proteins: hemoglobin, myoglobin, and cytochrome P450.

History

Main article: History of ferrous metallurgy

Wrought iron

File:Mars symbol.svg
The symbol for Mars has been used since antiquity to represent iron.
File:QtubIronPillar.JPG
The Delhi iron pillar is an example of the iron extraction and processing methodologies of India. The iron pillar at Delhi has withstood corrosion for the last 1600 years.

Iron objects of great age are much rarer than objects made of gold or silver due to the ease of corrosion of iron.[28] Beads made of meteoric iron in 3500 B.C. or earlier were found in Gerzah, Egypt by G. A. Wainwright.[29] 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.[29] Items that were likely made of iron by Egyptians date from 2500 to 3000 BC.[28] Iron had a distinct advantage over bronze in warfare implements. It was much harder and more durable than bronze, although susceptible to rust.

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 BC.[30] The Hittites appear to be the first to understand the production of iron from its ores and regard it highly in their society.[24] They 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.[30] 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.[30]

Artifacts from smelted iron occur in India from 1800–1200 BC,[31] and in the Levant from about 1500 BC (suggesting smelting in Anatolia or the Caucasus).[32][33]

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."[28] 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).[34] 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.[35]

According to Pliny the Elder, iron use was common in the Roman era. He reported in his "Natural History" chapter 39 that iron was used to help "construct houses, cleave rocks, and perform so many other useful offices of life" and that iron was used in "wars, murders, and robberies."[29]

Cast iron

Cast iron was first produced in China about 550 BC,[36] but was hardly in Europe until the medieval period.[37][38] 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.

Medieval blast furnaces were about 10 feet tall and made of fireproof brick; forced air was usually provided by hand-operated bellows.[38] Modern blast furnaces have grown much bigger.

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 wasn't implicated as the reason for the differences in properties of wrought iron, cast iron and steel until the 18th century.[30]

Steel

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 BC.[30] Then improved versions, Wootz steel by India and Damascus steel by China were developed around 300 B.C. and 500 A.D. respectively. These methods were specialized, and so steel did not become a major commodity until the 1850s.

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.(Citation needed)

Recent discoveries

Industrial production

See also: Iron ore

The production of iron or steel is a process containing two main stages, unless the desired product is cast iron. The first stage is to produce pig iron in a blast furnace. Alternatively, it may be directly reduced. The second is to make wrought iron or steel from pig iron by a further process.

File:Chinese Puddle and Blast Furnace.jpg
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.
File:Iron-Making.jpg
How iron was extracted in the 19th century

For a few limited purposes like electromagnet cores, pure iron is produced by electrolysis of a ferrous sulfate solution[24]

Blast furnace

Main article: Blast furnace

Ninety percent of all mining of metallic ores is for the extraction of iron[citation needed]. Industrially, iron production involves iron ores, principally hematite (nominally Fe2O3) and magnetite (Fe3O4) in a carbothermic reaction (reduction with carbon) in a blast furnace at temperatures of about 2000 °C. In a blast furnace, iron ore, carbon in the form of coke, and a flux such as limestone (which is used to remove silicon dioxide impurities in the ore which would otherwise clog the furnace with solid material) are fed into the top of the furnace, while a massive blast of heated air, about 4 tons per ton of iron,[38] is forced into the furnace at the bottom.

File:2005iron ore.PNG
Iron output in 2005

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

2 C + O2 → 2 CO

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

Fe2O3 + 3 CO → 2 Fe + 3 CO2

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

2 Fe2O3 + 3 C → 4 Fe + 3 CO2

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

CaCO3 → CaO + CO2

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

CaO + SiO2 → CaSiO3

The slag melts in the heat of the furnace. In the bottom of the furnace, the molten slag floats on top of the denser molten iron, and apertures in the side of the furnace are opened to run off the iron and the slag separately. The iron, once cooled, is called pig iron, while the slag can be used as a material in road construction or to improve mineral-poor soils for agriculture[38]

File:LightningVolt Iron Ore Pellets.jpg
This heap of iron ore pellets will be used in steel production.

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

Direct iron reduction

Since coke is becoming more regulated due to environmental concerns, alternative methods of processing iron have been developed. One of them is known as direct iron reduction.[38] It reduces iron ore to a powder substance called sponge iron, which is suitable for steelmaking. There are two main reactions that go on in the direct reduction process:

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

2 CH4 + O2 → 2 CO + 4 H2

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

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

Silica is removed by adding a flux, i.e. limestone, later.

Further processes

Main articles: Steelmaking and Ironworks

Pig iron is not pure iron, but has 4–5% carbon dissolved in it with small amounts of other impurities like sulfur, magnesium, phosphorus and manganese. As the carbon is the major impurity, the iron (pig iron) becomes brittle and hard. This form of iron, 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.

The hardness of the steel depends upon its carbon content: the higher the percentage of carbon, the greater the hardness and the lesser the malleability. The properties of the steel can also be changed by several methods.

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.

Steel may be hardened by cold working. The metal is bent or hammered into its final shape at a relatively cool temperature. Cold forging is the stamping of a piece of steel into shape by a heavy press. Wrenches are commonly made by cold forging. Cold rolling, which involves making a thinner but harder sheet, and cold drawing, which makes a thinner but stronger wire, are two other methods of cold working. To harden the steel, it is heated to red hot and then cooled by quenching it in the water. It becomes harder and more brittle. If it is too hardened, it is then heated to a required temperature and allowed to cool. The steel thus formed is less brittle.

Heat treatment is another way to harden steel. The steel is heated red hot, then cooled quickly. The iron carbide molecules are decomposed by the heat, but do not have time to reform. Since the free carbon atoms are stuck, it makes the steel much harder and stronger than before.[38]

Sometimes both toughness and hardness are desired. A process called case hardening may be used. Steel is heated to about 900 °C in a bed of charcoal and/or nitrogen. They diffuse into the steel, making the surface very hard. The surface cools quickly, but the inside cools slowly, making an extremely hard surface and a durable, resistant inner layer.

Iron may be passivated by dipping it into a concentrated nitric acid solution. This forms a protective layer of oxide on the metal, protecting it from further corrosion. When the metal is jarred, however, the layer is broken, allowing the metal to corrode again.[24]

Applications

Metallurgical

File:Ironattenuation.PNG
Photon mass attenuation coefficient for iron.

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

Commercially available iron is classified based on purity and the abundance of additives. Pig iron has 3.5–4.5% carbon[39] 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 from iron ore. Cast iron contains 2–4% carbon, 1–6% silicon, and small amounts of manganese. Contaminants present in pig iron that negatively affect material properties, such as sulfur and phosphorus, have been reduced to an acceptable level. It has a melting point in the range of 1420–1470 K, which is lower than either of its two main components, and makes it the first product to be melted when carbon and iron are heated together. Its mechanical properties vary greatly, dependent upon the form carbon takes in the alloy.

"White" cast irons contain their carbon in the form of cementite, or iron carbide. This hard, brittle compound dominates the mechanical properties of white cast irons, rendering them hard, but unresistant to shock. The broken surface of a white cast iron is full of fine facets of the broken carbide, a very pale, silvery, shiny material, hence the appellation.

In gray iron the carbon exists free as fine flakes of graphite, and also renders the material brittle due to the stress-raising nature of the sharp edged flakes of graphite. A newer variant of 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, vastly increasing the toughness and strength of the material.

Wrought iron contains less than 0.25% carbon.[39] It is a tough, malleable product, but not as fusible as pig iron. If honed to an edge, it loses it quickly. Wrought iron is characterized by the presence of fine fibers of slag entrapped in the metal. Wrought iron is more corrosion resistant than steel. It has been almost completely replaced by mild steel for traditional "wrought iron" products and blacksmithing.

Mild steel corrodes more readily that wrought iron, but is cheaper and more widely available. Carbon steel contains 2.0% carbon or less,[40] 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. Painting, galvanization, passivation, plastic coating and bluing are all used to protect iron from rust by excluding water and oxygen or by sacrificial protection.

Of 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.[41] Powdered iron in an acidic solvent was used in the Bechamp reduction the reduction of nitrobenzene to aniline.[42]

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.[43] It can also be dissolved in alcohol to form tincture of iron.[24] 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

File:Heme b.png
Structure of Heme b, in the protein additional ligand(s) would be attached to Fe.
Main article: Human iron metabolism

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, and catalase.[44]

Uptake and storage

In cells, iron storage is carefully regulated, "free" iron does not exist as such. A major component of this regulation is the protein transferrin, which binds iron absorbed from the duodenum and carries it in the blood to cells.[45] In animals, plants, and fungi, iron is often the metal ion incorporated into the heme complex. Heme is an essential component of cytochrome proteins, which mediate redox reactions, and of oxygen carrier proteins such as hemoglobin, myoglobin, and leghemoglobin. Inorganic iron also contributes to redox reactions in the iron-sulfur clusters of many enzymes, such as nitrogenase (involved in the synthesis of ammonia from nitrogen and hydrogen) and hydrogenase. Non-heme iron proteins include the enzymes methane monooxygenase (oxidizes methane to methanol), ribonucleotide reductase (reduces ribose to deoxyribose; DNA biosynthesis), hemerythrins (oxygen transport and fixation in Marine invertebrates) and purple acid phosphatase (hydrolysis of phosphate esters).

Iron distribution is heavily regulated in mammals, partly because iron has a high potential for biological toxicity.[46]

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.[47][48] [49]

Health and diet

Main article: Iron deficiency (medicine)

Iron is pervasive, but particularly rich sources of dietary iron include red meat, lentils, beans, poultry, fish, leaf vegetables, 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 (haem iron) is more easily absorbed than iron in vegetables.[50] Although most studies suggest that heme/hemoglobin from red meat has effects which may increase the likelihood of colorectal cancer,[51][52] there is still some controversy,[53] and even a few studies suggesting that there is not enough evidence to support such claims.[54]

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),[55] 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[56] and is also available for use as a common iron supplement. Often the amino acid chosen for this purpose is the cheapest and most common amino acid, glycine, leading to "iron glycinate" supplements.[57] The Recommended Dietary Allowance (RDA) for iron varies considerably based on age, gender, and source of dietary iron (heme-based iron has higher bioavailability).[58] Infants may require iron supplements if they are bottle-fed cow's milk.[59] Blood donors and pregnant women are at special risk of low iron levels and are often advised to supplement their iron intake.[60]

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 mostly by regulating uptake.[61] Regulation of iron uptake is impaired in some people as a result of a genetic defect that maps to the HLA-H gene region on chromosome 6. In these people, excessive iron intake can result in iron overload disorders, such as hemochromatosis. Many people have a genetic susceptibility to iron overload without realizing it or being aware of a family history of the problem. For this reason, it is advised that people do not take iron supplements unless they suffer from iron deficiency and have consulted a doctor. Hemochromatosis is estimated to cause disease in between 0.3 and 0.8% of Caucasians.[62]

MRI finds 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.[63]

Precautions

NFPA 704
1
1
Fire diamond for powdered iron metal
Main article: Iron poisoning

Large amounts of ingested iron can cause excessive levels of iron in the blood. High blood levels of free ferrous iron react with peroxides to produce free radicals, which are highly reactive and can damage DNA, proteins, lipids, and other cellular components. Thus, iron toxicity occurs when there is free iron in the cell, which generally occurs when iron levels exceed the capacity of transferrin to bind the iron. 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, which can cause significant adverse effects, including coma, metabolic acidosis, shock, liver failure, coagulopathy, adult respiratory distress syndrome, long-term organ damage, and even death.[64] Humans experience iron toxicity above 20 milligrams of iron for every kilogram of mass, and 60 milligrams per kilogram is considered a lethal dose.[65] 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.[65] The Dietary Reference Intake (DRI) lists the Tolerable Upper Intake Level (UL) for adults as 45 mg/day. For children under fourteen years old the UL is 40 mg/day.

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.[64][66]

See also

References

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Books

  • 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.
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External links

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

af:Yster als:Eisen ar:حديد an:Fierro arc:ܦܪܙܠܐ ast:Fierro gn:Kuarepoti az:Dəmir bn:লোহা zh-min-nan:Fe (goân-sò͘) be:Жалеза be-x-old:Жалеза bs:Željezo br:Houarn bg:Желязо ca:Ferro cv:Тимĕр cs:Železo co:Ferru cy:Haearn da:Jern de:Eisen dv:ދަގަނޑު nv:Béésh (Fe) et:Raud el:Σίδηρος myv:Кшни es:Hierro eo:Fero ext:Hierru eu:Burdina fa:آهن fr:Fer fur:Fier ga:Iarann gv:Yiarn gd:Iarann gl:Ferro gan:鐵 gu:લોખંડ hak:Thiet xal:Төмр ko:철 hy:Երկաթ hi:लोहा hr:Željezo io:Fero id:Besi ia:Ferro is:Járn it:Ferro he:ברזל jv:Wesi kn:ಕಬ್ಬಿಣ ka:რკინა kk:Темір sw:Chuma kv:Кӧрт kg:Kibende ht:Fè ku:Hesin la:Ferrum lv:Dzelzs lb:Eisen lt:Geležis lij:Færo (elemento) li:Iezer ln:Ebendé jbo:tirse hu:Vas mk:Железо mg:Vy ml:ഇരുമ്പ് mt:Ħadid mi:Rino mr:लोखंड ms:Besi mn:Төмөр (химийн элемент) nah:Tlīltic tepoztli nl:IJzer (element) ja:鉄 frr:Stälj no:Jern nn:Jern nrm:Fé oc:Fèrre uz:Temir pa:ਲੋਹਾ pnb:لوآ nds:Iesen pl:Żelazo pt:Ferro ksh:Eisen ro:Fier qu:Khillay ru:Железо sah:Тимир sa:लोहम् sco:Airn stq:Iersen sq:Hekuri scn:Ferru simple:Iron sk:Železo sl:Železo so:Bir sr:Гвожђе (хемијски елемент) sh:Željezo fi:Rauta sv:Järn tl:Bakal ta:இரும்பு te:ఇనుము th:เหล็ก tg:Оҳан tr:Demir uk:Залізо ur:لوہا ug:تۆمۈر za:Diet vec:Fero vi:Sắt fiu-vro:Raud vls:Yzer (element) war:Putháw yi:אייזן yo:Iron zh-yue:鐵 bat-smg:Gelžis

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  1. ↑ 1.0 1.1 Kohl, Walter H. (1995). Handbook of materials and techniques for vacuum devices. Springer. pp. 164–167. ISBN 1563963876. 
  2. ↑ 2.0 2.1 Kuhn, Howard and Medlin, Dana (prepared under the direction of the ASM International Handbook Committee), ed. (2000). ASM Handbook – Mechanical Testing and Evaluation (PDF). 8. ASM International. p. 275. ISBN 0871703890. 
  3. ↑ "Hardness Conversion Chart". Maryland Metrics. Retrieved 2010-05-23. 
  4. ↑ Raghavan, V. (2004). Materials Science and Engineering. PHI Learning Pvt. Ltd. p. 218. ISBN 8120324552. 
  5. ↑ Takaji, Kusakawa; Toshikatsu, Otani. "Properties of Various Pure Irons : Study on pure iron I". Tetsu-to-Hagane. 50 (1): 42–47. 
  6. ↑ 6.0 6.1 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 9780871707482. 
  7. ↑ Martin, John Wilson (2007). Concise encyclopedia of the structure of materials. Elsevier. p. 183. ISBN 0080451276. 
  8. ↑ Georges, Audi (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A. Atomic Mass Data Center. 729: 3–128. doi:10.1016/j.nuclphysa.2003.11.001. 
  9. ↑ 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. 
  10. ↑ "Iron and Nickel Abundances in H~II Regions and Supernova Remnants". June 14, 1995. Retrieved 2010-05-23. .
  11. ↑ "Iron: geological information". WebElements. Retrieved 2010-05-23. .
  12. ↑ John W. Morgan*and Edward Anders (1980). "Chemical composition of Earth, Venus, and Mercury". Proc. Nat. Acad. Sci. 77 (12): 6973 –6977. doi:10.1073/pnas.77.12.6973. 
  13. ↑ Lyons, T. W.; Reinhard, CT (2009). "Early Earth: Oxygen for heavy-metal fans". Nature. 461 (7261): 179–181. doi:10.1038/461179a. PMID 19741692. 
  14. ↑ Cloud, P. (1973). "Paleoecological Significance of the Banded Iron-Formation". Economic Geology. 68: 1135–1143. doi:10.2113/gsecongeo.68.7.1135. 
  15. ↑ Emiliani, Cesare (1992). "Planet earth: cosmology, geology, and the evolution of life and environment". Cambridge University Press: 152. ISBN 9780521409490.  |chapter= ignored (help)
  16. ↑ 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. doi:10.1007/s10751-007-9508-5. 
  17. ↑ 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. 
  18. ↑ 18.0 18.1 18.2 18.3 18.4 18.5 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 3110075113. 
  19. ↑ 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 9780306416477. 
  20. ↑ 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 9781900747073. 
  21. ↑ Gmelin, Leopold (1852). "Mercury and Iron". Hand-book of chemistry. 6. Cavendish Society. pp. 128–129. 
  22. ↑ "Questions and Answers on Prussian Blue". Retrieved 2009-06-06. 
  23. ↑ 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. 
  24. ↑ 24.0 24.1 24.2 24.3 24.4 "Iron." Microsoft® Student 2009 [DVD]. Redmond, WA: Microsoft Corporation, 2008.
  25. ↑ Greenwood, Norman N.; Earnshaw, A. (1984), Chemistry of the Elements, Oxford: Pergamon, pp. 1282–86, ISBN 0-08-022057-6 .
  26. ↑ Pierre Laszlo and Roald Hoffmann (2000) "Ferrocene: Ironclad History of Rashomon Tale?," Angewandte Chemie (International Edition), vol. 39, no. 1, pages 123-124.
  27. ↑ Federman Neto, Alberto; Pelegrino, Alessandra Caramori; Darin, Vitor Andre (2004). "Ferrocene: 50 Years of Transition Metal Organometallic Chemistry — From Organic and Inorganic to Supramolecular Chemistry". ChemInform. 35. doi:10.1002/chin.200443242. 
  28. ↑ 28.0 28.1 28.2 Weeks 1968, p. 29.
  29. ↑ 29.0 29.1 29.2 Weeks 1968, p. 31.
  30. ↑ 30.0 30.1 30.2 30.3 30.4 Weeks 1968, p. 32.
  31. ↑ 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 2010-05-23. 
  32. ↑ 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. 
  33. ↑ Muhly, James D. (2003). "Metalworking/Mining in the Levant". In Lake, Richard Winona. Near Eastern Archaeology IN: Eisenbrauns. 180. pp. 174–183. 
  34. ↑ Weeks 1968, pp. 29–30.
  35. ↑ Weeks 1968, p. 30.
  36. ↑ 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. 
  37. ↑ Giannichedda, Enrico (2007): "Metal production in Late Antiquity", in Technology in Transition AD 300-650 L. Lavan E.Zanini & A. Sarantis Brill, eds., Leiden; p. 200
  38. ↑ 38.0 38.1 38.2 38.3 38.4 38.5 Biddle, Verne; Parker, Gregory. Chemistry, Precision and Design. A Beka Book, Inc. 
  39. ↑ 39.0 39.1 Camp, James McIntyre; Francis, Charles Blaine (1920). The Making, Shaping and Treating of Steel. Pittsburgh: Carnegie Steel Company. pp. 173–174. 
  40. ↑ "Classification of Carbon and Low-Alloy Steels". Retrieved 2008-01-05. 
  41. ↑ Kolasinski, Kurt W. (2002). "Where are Heterogenous Reactions Important". Surface science: foundations of catalysis and nanoscience. John Wiley and Sons. pp. 15–16. ISBN 9780471492443. 
  42. ↑ 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 9780824724818. 
  43. ↑ Wildermuth, Egon; Stark, Hans; Friedrich, Gabriele; Ebenhöch, Franz Ludwig; Kühborth, Brigitte; Silver, Jack; Rituper, Rafael (2000). "Iron Compounds". doi:10.1002/14356007.a14_591. 
  44. ↑ Lippard, S. J.; Berg, J. M. (1994). Principles of Bioinorganic Chemistry. Mill Valley: University Science Books. ISBN 0-935702-73-3. 
  45. ↑ Rouault, Tracey A. (2003). "How Mammals Acquire and Distribute Iron Needed for Oxygen-Based Metabolism". PLoS Biology. 1 (1): e9. doi:10.1371/journal.pbio.0000079. PMC 212690 Freely accessible. PMID 14551907. 
  46. ↑ Nanami, M.; Ookawara, T; Otaki, Y; Ito, K; Moriguchi, R; Miyagawa, K; Hasuike, Y; Izumi, M; Eguchi, H (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. 
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  49. ↑ Boukhalfa, Hakim; Crumbliss, Alvin L. (2002). "Chemical aspects of siderophore mediated iron transport". BioMetals. 15 (4): 325. doi:10.1023/A:1020218608266. PMID 12405526. 
  50. ↑ Food Standards Agency - Eat well, be well - Iron deficiency
  51. ↑ 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. PMID 10582688. 
  52. ↑ 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. 
  53. ↑ 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. PMID 11352852. 
  54. ↑ "Eating Red Meat Will Not Increase Colorectal Cancer Risk, Study Suggests". ScienceDaily. June 13, 2007. Retrieved 2010-05-23. 
  55. ↑ [1]
  56. ↑ 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. 
  57. ↑ Ashmead, H. DeWayne (1989). Conversations on Chelation and Mineral Nutrition. Keats Publishing. ISBN 0-87983-501-X. 
  58. ↑ "Dietary Reference Intakes: Elements" (PDF). The National Academies. 2001. Retrieved 2008-05-21. 
  59. ↑ "Iron Deficiency Anemia" (web page). MediResource. Retrieved 2008-12-17. 
  60. ↑ 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. 
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  63. ↑ 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. 
  64. ↑ 64.0 64.1 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. 
  65. ↑ 65.0 65.1 "Toxicity, Iron". Medscape. Retrieved 2010-05-23. 
  66. ↑ 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. 
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