Iron

26 manganese ← iron → cobalt
-
↑
Fe
↓
Ru
Iron in the periodic table of the elements
Periodic table - Extended periodic table
General
Name, symbol, number iron, Fe, 26
Element category transition metal
Group, period, block 8, 4, d
Appearance lustrous metallic
with a grayish tinge
Fe,26.jpg
Standard atomic weight 55.845(2) g·mol−1
Electron configuration [Ar] 3d6 4s2
Electrons per shell 2, 8, 14, 2
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
Crystal structure body-centered cubic
a=286.65 pm;
face-centered cubic
between 1185–1667 K
Oxidation states 6, 5 [1], 4, 3, 2, 1 [2]
(amphoteric oxide)
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 140 pm
Atomic radius (calc.) 156 pm
Covalent radius 125 pm
Miscellaneous
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 4.0
Vickers hardness 608 MPa
Brinell hardness 490 MPa
CAS registry number 7439-89-6
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% 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 1.5×106 y β- 3.978 60Co
References

Iron (pronounced /ˈаɪɚn/) is a chemical element with the symbol Fe (Latin: ferrum) and atomic number 26. Iron is a group 8 and period 4 element. Iron is lustrous and silvery in color. It is soft, about 80 Brinell,[1] relative to steel, which is about 140 Brinell.[2] It is one of the few ferromagnetic elements.

Iron is the heaviest element produced by stellar nucleosynthesis; heavier elements require a red giant or supernova for their formation. Iron and nickel are the most abundant metals in iron meteorites and in the dense metal cores of planets such as Earth. Iron and iron alloys are also the most common source of ferromagnetic materials in everyday use.

Contents

Occurrence

See also: Category:Iron minerals

Iron is the sixth most abundant element in the universe, formed as the final act of nucleosynthesis by carbon burning in massive stars. While it makes up about 5% of the Earth's crust, the earth's core is 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.[3] Most of the iron in the crust is found combined with oxygen as iron oxide minerals such as hematite and magnetite. About 1 in 20 meteorites consist of the unique iron-nickel minerals taenite (35–80% iron) and kamacite (90–95% iron). Although rare, meteorites are the major form of natural metallic iron on the earth's surface.

The reason for Mars' red colour is thought to be an iron-oxide-rich soil.

The red appearance of this water is due to iron in the rocks.

Characteristics

Iron is a metal extracted mainly from the iron ore hematite. It oxidises readily in air and water to form Fe2O3 and is rarely found as a free element. In order to obtain elemental iron, oxygen and other impurities must be removed by chemical reduction. The properties of iron can be modified by alloying it with various other metals and some non-metals, notably carbon and silicon to form steels.

Nuclei of iron atoms have some of the highest binding energies per nucleon, surpassed only by the nickel isotope 62Ni. The universally most abundant of the highly stable nuclides is, however, 56Fe. 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 favoured, and iron abundance on Earth greatly favors iron over nickel, and also presumably in supernova element production.[4]

Iron (as Fe2+, ferrous ion) is a necessary trace element used by almost all living organisms. The only exceptions are several organisms that live in iron-poor environments and have evolved to use different elements in their metabolic processes, such as manganese instead of iron for catalysis, or 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.

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 alpha, gamma, and delta.

As molten iron cools down it crystallises at 1538 °C into its delta allotrope, which has a body-centred cubic (BCC) crystal structure. As it cools further its crystal structure changes to face-centred cubic (FCC) at 1394 °C, when it is known as gamma-iron, or austenite. At 912 °C the crystal structure again becomes BCC as alpha-iron is formed, and at 770 °C (the Curie point, Tc) the 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 unmagnetised iron, all the electronic spins of the atoms within one domain are in the same direction; however, in neighbouring domains they point in various directions and thus cancel out. In magnetised iron, the electronic spins of all the domains are all aligned, so that the magnetic effects of neighbouring domains reinforce each other. Although each domain contains billions of atoms, they are very small, about one thousandth of a centimetre across.

Iron is of most 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).

Above 912 °C and up to 1401 °C alpha iron undergoes a phase transition from body-centred cubic to the face-centred cubic configuration of gamma 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.

Applications

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

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 some techniques used to protect iron from rust by excluding water and oxygen or by sacrificial protection.

Iron compounds

Iron chloride hexahydrate

History

Main article: History of ferrous metallurgy
The symbol for Mars has been used since ancient times to represent iron.
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.
The puddling 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 1637 by Song Yingxing.

The first iron used by mankind during prehistory came from meteors. The smelting of iron in bloomeries probably began in Anatolia, India or the Caucasus in the second millennium BC or the later part of the preceding one. Cast iron was first produced in China about 550 BC, but not in Europe until the medieval period. 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.

Steel (with a smaller carbon content than pig iron but more than wrought iron) was first produced in antiquity. 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 and other 19th century and later processes have led to wrought iron no longer being produced.

Production of iron from iron ore

See also: Iron ore

The production of iron or steel is a process unless the desired final product is cast iron. The first stage is to produce pig iron in a blast furnace. The second is to make wrought iron or steel from pig iron by a further process.

Blast furnace

Main article: Blast furnace
Iron output in 2005

Ninety percent of all mining of metallic ores is for the extraction of iron. Industrially, iron is produced starting from iron ores, principally haematite (nominally Fe2O3) and magnetite (Fe3O4) by 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 impurities in the ore which would otherwise clog the furnace with solid material) are fed into the top of the furnace, while a blast of heated air is forced into the furnace at the bottom.

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:

3 CO + Fe2O3 → 2 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 (quicklime):

CaCO3 → CaO + CO2

Then calcium oxide combines with silicon dioxide to form a 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.

How Iron was extracted in the 19th century
This heap of iron ore pellets will be used in steel production.

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

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 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 oxidise 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 proportion of carbon, the greater the hardness and the lesser the ductility. The properties of the steel can also be changed by tempering it. 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. This steel is then heated to a required temperature and allowed to cool. The steel thus formed is less brittle.

Isotopes

Main article: Isotopes of iron

Naturally occurring iron consists of four isotopes: 5.845% of radioactive 54Fe (half-life: >3.1×1022 years), 91.754% of stable 56Fe, 2.119% of stable 57Fe and 0.282% of stable 58Fe. 60Fe is an extinct radionuclide of long half-life (1.5 million years).

Much of the past work on measuring the isotopic composition of Fe has centered 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.[10]

The isotope 56Fe is of particular interest to nuclear scientists. A common misconception is that this isotope represents the most stable nucleus possible, and that it thus would be impossible to perform fission or fusion on 56Fe and still liberate energy. This is not true, as both 62Ni and 58Fe are more stable, being the most stable nuclei. However, since 56Fe is much more easily produced from lighter nuclei in nuclear reactions, it is the endpoint of fusion chains inside extremely massive stars and is therefore common in the universe, relative to other metals.

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).

Iron in organic synthesis

The use of iron metal filings in organic synthesis is mainly for the reduction of nitro compounds.[11] Additionally, iron has been used for desulfurizations,[12] reduction of aldehydes,[13] and the deoxygenation of amine oxides.[14]

Iron in biology

Structure of Heme b
Main article: Human iron metabolism

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. Iron deficiency can lead to iron deficiency anemia.

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.[15]

Nutrition and dietary sources

Good sources of dietary iron include red meat, fish, poultry, lentils, beans, leaf vegetables, tofu, chickpeas, black-eyed peas, fortified bread, and fortified breakfast cereals. Iron in low amounts is found in molasses, teff and farina. Iron in meat is more easily absorbed than iron in vegetables (haem iron),[16] but heme/hemoglobin from red meat has effects which may increase the likelihood of colorectal cancer.[17][18]

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[19] 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.[20] The RDA for iron varies considerably based on age, gender, and source of dietary iron (heme-based iron has higher bioavailability).[21] Infants may require iron supplements if they are not breast-fed. Blood donors and pregnant women 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 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.[22] However, large amounts of ingested iron can cause excessive levels of iron in the blood because high iron levels can damage the cells of the gastrointestinal tract, preventing 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.[23] 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.[23] 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 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.[24]

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.

See also

Bibliography

References

  1. ↑ Structure of plain steel, http://www.key-to-steel.com/articles/art3.htm, retrieved on 2008-10-21 .
  2. ↑ IDE 120 - Hardness Testing, http://www.google.com/url?sa=t&source=web&ct=res&cd=1&url=http%3A%2F%2Fweb.mst.edu%2F~ide120%2Fextra%2Fdata_sheets%2Fhardness.doc&ei=z27-SNjwH4i6NPr7yNkC&usg=AFQjCNGopQh_My1yK6R1ej0mhbs9Mlb9ow&sig2=w3avAIETD_98MQ7A8f3KnQ, retrieved on 2008-10-21 .
  3. ↑ Iron: geological information, http://www.webelements.com/iron/geology.html, retrieved on 2008-05-21 .
  4. ↑ Iron and Nickel Abundances in H~II Regions and Supernova Remnants, June 14, 1995, http://www.aas.org/publications/baas/v27n2/aas186/abs/S3707.html, retrieved on 2008-05-21 .
  5. ↑ 5.0 5.1 Camp, James McIntyre; Francis, Charles Blaine (1920). The Making, Shaping and Treating of Steel. Pittsburgh: Carnegie Steel Company. pp. 173–174. http://books.google.com/books?id=P9MxAAAAMAAJ. 
  6. ↑ Classification of Carbon and Low-Alloy Steels, http://www.key-to-steel.com/Articles/Art62.htm, retrieved on 2008-01-05 
  7. ↑ Vivian Marx (2002). "The Little Plankton That Could…Maybe". Scientific American. http://www.sciam.com/article.cfm?articleID=000A5750-8AC2-1D9C-815A809EC5880000. 
  8. ↑ Melinda Ferguson, David Labiak, Andrew Madden, Joseph Peltier. "The Effect of Iron on Plankton Use of CO2". CEM 181H. Retrieved on 2007-05-05.
  9. ↑ Dopyera, Caroline (October, 1996). "The Iron Hypothesis". EARTH. Retrieved on 2007-05-05.
  10. ↑ Dauphas, N. & Rouxel, O. 2006. Mass spectrometry and natural variations of iron isotopes. Mass Spectrometry Reviews, 25, 515-550
  11. ↑ Fox, B. A.; Threlfall, T. L. Organic Syntheses, Coll. Vol. 5, p.346 (1973); Vol. 44, p.34 (1964). (Article)
  12. ↑ Blomquist, A. T.; Dinguid, L. I. J. Org. Chem. 1947, 12, 718 & 723.
  13. ↑ Clarke, H. T.; Dreger, E. E. Org. Syn., Coll. Vol. 1, p.304 (1941); Vol. 6, p.52 (1926). (Article).
  14. ↑ den Hertog, J.; Overhoff, J. Recl. Trav. Chim. Pays-Bas 1950, 69, 468.
  15. ↑ Tracey A. Rouault. "How Mammals Acquire and Distribute Iron Needed for Oxygen-Based Metabolism". Retrieved on 2006-06-19.
  16. ↑ Food Standards Agency - Eat well, be well - Iron deficiency
  17. ↑ Sesink AL, Termont DS, Kleibeuker JH, Van der Meer R (1999). "Red meat and colon cancer: the cytotoxic and hyperproliferative effects of dietary heme". Cancer Research 59 (22): 5704–9. PMID 10582688. http://cancerres.aacrjournals.org/cgi/pmidlookup?view=long&pmid=10582688. 
  18. ↑ Glei M, Klenow S, Sauer J, Wegewitz U, Richter K, Pool-Zobel BL (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): 162–71. doi:10.1016/j.mrfmmm.2005.08.006. PMID 16226281. 
  19. ↑ Pineda O, Ashmead HD (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. 
  20. ↑ Ashmead, H. DeWayne (1989). 'Conversations on Chelation and Mineral Nutrition'. Keats Publishing. ISBN 0-87983-501-X. 
  21. ↑ "Dietary Reference Intakes: Elements" (PDF). The National Academies (2001). Retrieved on 2008-05-21.
  22. ↑ Kumar, Vinay; Abbas, Abul K; Fausto, Nelson (2005). "Anemia". Robbins and Cotran: Pathologic Basis of Disease, 7th edition. Elsevier Saunders. Retrieved on 2008-03-14.
  23. ↑ 23.0 23.1 "Toxicity, Iron". Emedicine. Retrieved on 2006-06-19.
  24. ↑ Durupt S, Durieu I, Nove-Josserand R, et al: [Hereditary hemochromatosis]. Rev Med Interne 2000 Nov; 21(11): 961-71[Medline].

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