Iron

<tr><td>Chemical series</td><td>transition metals</td></tr> <tr><td>Appearance</td><td>lustrous metallic
with a grayish tinge
<tr><td>Atomic radius</td><td>140 pm</td></tr><tr><td>Atomic radius (calc.)</td><td>156 pm</td></tr><tr><td>Covalent radius</td><td>125 pm</td></tr> <tr><td rowspan="2" valign="top">Magnetic ordering</td><td>ferromagnetic</td></tr><tr><td>1043 K</td></tr><tr><td>Electrical resistivity</td><td>(20 °C) 96.1 nΩ·m</td></tr><tr><td>Thermal conductivity</td><td>(300 K) 80.4 W·m⁻¹·K⁻¹</td></tr><tr><td>Thermal expansion</td><td>(25 °C) 11.8 µm·m⁻¹·K⁻¹</td></tr><tr><td>Speed of sound (thin rod)</td><td>(r.t.) (electrolytic)
5120 m·s⁻¹</td></tr><tr><td>Young's modulus</td><td>211 GPa</td></tr><tr><td>Shear modulus</td><td>82 GPa</td></tr><tr><td>Bulk modulus</td><td>170 GPa</td></tr><tr><td>Poisson ratio</td><td>0.29</td></tr><tr><td>Mohs hardness</td><td>4.0</td></tr><tr><td>Vickers hardness</td><td>608 MPa</td></tr><tr><td>Brinell hardness</td><td>490 MPa</td></tr><tr><td>CAS registry number</td><td>7439-89-6</td></tr>

26 manganese ← iron → cobalt - ↑ Fe ↓ Ru Periodic table - Extended periodic table
General
Name, symbol, number	iron, Fe, 26
Group, period, block	8, 4, d
Standard atomic weight	55.845(2) g·mol−1
Electron configuration	[Ar] 4s2 3d6
Electrons per shell	2, 8, 14, 2
Physical properties<tr><td>Phase</td><td>solid</td></tr><tr><td>Density (near r.t.)</td><td>7.86 g·cm−3</td></tr><tr><td>Liquid density at m.p.</td><td>6.98 g·cm−3</td></tr><tr><td>Melting point</td><td>1811 K (1538 °C, 2800 °F)</td></tr><tr><td>Boiling point</td><td>3134 K (2861 °C, 5182 °F)</td></tr><tr><td>Heat of fusion</td><td>13.81 kJ·mol−1</td></tr><tr><td>Heat of vaporization</td><td>340 kJ·mol−1</td></tr><tr><td>Heat capacity</td><td>(25 °C) 25.10 J·mol−1·K−1</td></tr>
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K 1728 1890 2091 2346 2679 3132
Atomic properties <tr><td>Crystal structure</td><td>body-centered cubic a=286.65 pm; face-centered cubic between 1185–1667 K</td></tr><tr><td>Oxidation states</td><td>2, 3, 4, 6 (amphoteric oxide)</td></tr><tr><td>Electronegativity</td><td>1.83 (Pauling scale)</td></tr>
Ionization energies (more)	1st: 762.5 kJ·mol−1
	2nd: 1561.9 kJ·mol−1
	3rd: 2957 kJ·mol−1
Miscellaneous
Selected 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% Fe is stable with 30 neutrons 57Fe 2.2% Fe is stable with 31 neutrons 58Fe 0.28% Fe is stable with 32 neutrons 59Fe syn 44.503 d β- 1.565 59Co 60Fe syn 1.5×106 y β- 3.978 60Co
References

Editor-In-Chief: C. Michael Gibson, M.S., M.D. [1]

Iron (pronounced /ˈaɪɚn/) is a chemical element with the symbol Fe and atomic number 26. Iron is a group 8 and period 4 element. Iron is a lustrous, silvery soft metal. It is one of the few ferromagnetic elements.

Characteristics

Iron is a metal extracted mainly from the iron ore hematite. It oxidises readily in air and water and is rarely found as a free element. In order to obtain elemental iron, oxygen and other impurities must be removed by chemical reduction. Iron is the main constituent of steel, and it is used in the production of alloys or solid solutions of various metals, as well as some non-metals, particularly carbon.

Iron (as Fe²⁺, ferrous ion) is a necessary trace element used by almost all living organisms, the only exceptions are a few prokaryotic organisms which live in iron-poor conditions (such as the lactobacilli in iron-poor milk) which use manganese for catalysis instead as well as organisms which use hemocyanin instead of hemoglobin. Iron-containing enzymes, usually containing heme prosthetic groups, participate in catalysis of oxidation reactions in biology, and in transport of a number of soluble gases. See hemoglobin, cytochrome, and catalase.

Applications

Iron is the most used of all the metals, comprising 95% of all the metal tonnage produced worldwide. Its combination of low cost and high strength make it indispensable, especially in applications like automobiles, the hulls of large ships, and structural components for buildings. Steel is the best known alloy of iron.

Iron compounds

• Iron(III) phosphate (FePO₄) is used in fertilizer and as an additive in human and animal food.

• Iron(II) gluconate (Fe(C₆H₁₁O₇)₂) is used as a dietary supplement in iron pills.

• Iron(II) sulfate (FeSO₄) is used in water purification and sewage treatment systems, as a catalyst in the production of ammonia, as an ingredient in fertilizer and herbicide, as an additive in animal feed, in wood preservative and as an additive to flour to increase iron levels.

Iron in organic synthesis

The usage of iron metal filings in organic synthesis is mainly for the reduction of nitro compounds. Additionally, iron has been used for desulfurizations, reduction of aldehydes, and the deoxygenation of amine oxides.

Iron in biology

Error creating thumbnail: /home/webapps/wikidoc/mediawiki-1.19.2/bin/ulimit4.sh: line 4: r: command not found

Structure of Heme b

Iron is essential to nearly all known organisms. In cells, iron is generally stored in the centre of metalloproteins, because "free" iron -- which binds non-specifically to many cellular components -- can catalyse production of toxic free radicals.

In animals, plants, and fungi, iron is often incorporated into the heme complex. Heme is an essential component of cytochrome proteins, which mediate redox reactions, and of oxygen carrier proteins such as hemoglobin, myoglobin, and leghemoglobin. 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. Iron distribution is also regulated because many bacteria require iron, so restricting its availability to bacteria (generally by sequestering it inside cells) can help to prevent or limit infections. This is probably the reason for the relatively low amounts of iron in mammalian milk. 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.

Nutrition and dietary sources

Good sources of dietary iron include red meat, fish, poultry, lentils, beans, leaf vegetables, tofu, chickpeas, black-eyed peas, potatoes with skin, bread made from completely whole-grain flour, molasses, teff and farina. Iron in meat is more easily absorbed than iron in vegetables.

Iron provided by dietary supplements is often found as iron (II) fumarate, although iron sulfate is cheaper and is absorbed equally well. Elemental iron, despite being absorbed to a much smaller extent (stomach acid is sufficient to convert some of it to ferrous iron), is often added to foods such as breakfast cereals or "enriched" wheat flour (where it is listed as "reduced iron" in the list of ingredients). Iron is most available to the body when chelated to amino acids - iron in this form is ten to fifteen times more bioavailable than any other, 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. The RDA for iron varies considerably based on age, gender, and source of dietary iron (heme-based iron has higher bioavailability). Infants will require iron supplements if they are not breast-fed. Blood donors are at special risk of low iron levels and are often advised to supplement their iron intake.

Regulation of iron uptake

Excessive iron can be toxic, because free ferrous iron reacts 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.

Iron uptake is tightly regulated by the human body, which has no physiological means of excreting iron, so controls iron levels solely by regulating uptake. Although uptake is regulated, large amounts of ingested iron can cause excessive levels of iron in the blood, because high iron levels can cause damage to the cells of the gastrointestinal tract that prevents them from regulating iron absorption. High blood concentrations of iron damage cells in the heart, liver and elsewhere, which can cause serious problems, including long-term organ damage and even death.

Humans experience iron toxicity above 20 milligrams of iron for every kilogram of mass, and 60 milligrams per kilogram is a lethal dose. Over-consumption 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. The 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.

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 should 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.

The medical management of iron toxicity is complex, and can include use of a specific chelating agent called deferoxamine to bind and expel excess iron from the body.

Bibliography

• Los Alamos National Laboratory — Iron
• 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.
• Crystal structure of iron

References

• Doulias PT, Christoforidis S, Brunk UT, Galaris D. Endosomal and lysosomal effects of desferrioxamine: protection of HeLa cells from hydrogen peroxide-induced DNA damage and induction of cell-cycle arrest. Free Radic Biol Med. 2003;35:719-28.

External links

• WebElements.com – Iron
• It's Elemental – Iron
• The Most Tightly Bound Nuclei

<span id="interwiki-af-fa" /> <span id="interwiki-eo-fa" /> <span id="interwiki-pt-fa" /> <span id="interwiki-ru-fa" />