Iron

26 manganeseironcobalt
-

Fe

Ru
Periodic Table - Extended Periodic Table
General
Name, Symbol, Number iron, Fe, 26
Chemical series transition metals
Group, Period, Block 8, 4, d
Appearance lustrous metallic
with a grayish tinge
Atomic mass 55.845(2) g/mol
Electron configuration [Ar] 3d6 4s2
Electrons per shell 2, 8, 14, 2
Physical properties
Phase solid
Density (near r.t.) 7.86 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
(2861 °C, 5182 °F)
Heat of fusion 13.81 kJ·mol−1
Heat of vaporization 340 kJ·mol−1
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;[1]
face-centered cubic
between 1185–1667 K
Oxidation states 2, 3, 4, 6
(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
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% 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

Iron (IPA: /ˈaɪə(ɹ)n/) is a chemical element with the symbol Fe (Latin: ferrum) and atomic number 26. Iron is a group 8 and period 4 metal. Iron and nickel are notable for being the final elements produced by stellar nucleosynthesis, and thus the heaviest elements which do not require a supernova or similarly cataclysmic event for formation. Iron and nickel are therefore the most abundant metals in metallic meteorites and in the dense-metal cores of planets such as Earth.

Contents

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Characteristics

Iron is believed to be the tenth most abundant element in the universe. Iron makes up 5% of the Earth's crust and is second in abundance to aluminium among the metals and fourth in abundance among the elements. Iron is also the most abundant element by mass, making up 35% of the mass of the Earth as a whole. The concentration of iron in the various layers of the Earth ranges from very high at the inner core to only a few percent in the outer crust.

Iron is a metal extracted from iron ore, and is almost never found in the free elemental state. In order to obtain elemental iron, the impurities must be removed by chemical reduction. Iron is the main component 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. The many iron-carbon allotropes, which have very different properties, are discussed in the article on steel.

Nuclei of iron have some of the highest binding energies per nucleon, surpassed only by the nickel isotope 62Ni. The universally most abundant of the highly stable nucleides is, however, 56Fe. This is formed by nuclear fusion in the stars. Although a further tiny energy gain could be extracted by synthesizing 62Ni, conditions in stars are not right for this process to be favoured, and iron abundance on Earth greatly favors iron over nickel, and also presumably in supernova element production. [citation needed] When a very large star contracts at the end of its life, internal pressure and temperature rise, allowing the star to produce progressively heavier elements, despite these being less stable than the elements around mass number 60, known as the "iron group". This leads to a supernova. Some cosmological models with an open universe predict that there will be a phase where as a result of slow fusion and fission reactions, everything will become iron.[citation needed]

Iron (as Fe2+, ferrous ion) is a necessary trace element used by all known living organisms. 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.

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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, and some of the forms that iron can take include:

The main drawback to iron and steel is that pure iron, and most of its alloys, suffer badly from rust if not protected in some way. Painting, galvanization, plastic coating and bluing are some techniques used to protect iron from rust by excluding water and oxygen or by sacrificial protection.

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Iron Compounds

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History

The first iron used by mankind, far back in prehistory, came from meteors. The smelting of iron in bloomeries probably began in Anatolia or the Caucasus in the second millennium BC or the latter 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.

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Occurrence

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

Iron is one of the most common elements on Earth, making up about 5% of the Earth's crust. Most of this iron is found in various iron oxides, such as the minerals hematite, magnetite, and taconite. The earth's core is believed to consist largely of a metallic iron-nickel alloy. About 5% of the meteorites similarly consist of iron-nickel alloy. Although rare, these are the major form of natural metallic iron on the earth's surface.

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

See also iron minerals.

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Production of iron from iron ore

It has been suggested that this article or section be merged into Blast furnace. (Discuss)
How Iron was extracted in the 19th century
How Iron was extracted in the 19th century
This heap of iron ore pellets will be used in steel production.
This heap of iron ore pellets will be used in steel production.

Industrially, iron is produced starting from iron ores, principally hematite (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 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:

6 C + 3 O2 → 6 CO

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

6 CO + 2 Fe2O3 → 4 Fe + 6 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 (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, which silicon dioxide would not have. In the bottom of the furnace, the molten slag floats on top of the more dense 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.

Pig iron is not pure iron, but has 4-5% carbon dissloved in it. This is subsequently reduced to steel or commercially pure iron, known as wrought iron, using other furnaces or converters.

Approximately 1100Mt (million tons) of iron ore was produced in the world in 2000, with a gross market value of approximately 25 billion US dollars. While ore production occurs in 48 countries, the five largest producers were China, Brazil, Australia, Russia and India, accounting for 70% of world iron ore production. The 1100Mt of iron ore was used to produce approximately 572Mt of pig iron.

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Isotopes

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.[2]

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.

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).

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Iron in biology

Structure of Heme b
Structure of Heme b

Iron is essential to nearly all known organisms. It is mostly stably incorporated in the inside of metalloproteins, because in exposed or in free form it causes production of free radicals that are generally toxic to cells. To say that iron is free doesn't mean that it is free floating in the bodily fluids. Iron binds avidly to virtually all biomolecules so it will adhere nonspecifically to cell membranes, nucleic acids, proteins etc.

Many animals incorporate iron into the heme complex, an essential component of cytochromes, which are proteins involved in redox reactions (including but not limited to cellular respiration), and of oxygen carrying proteins hemoglobin and myoglobin. Inorganic iron involved in redox reactions is also found in the iron-sulfur clusters of many enzymes, such as nitrogenase (involved in the synthesis of ammonia from nitrogen and hydrogen) and hydrogenase. A class of non-heme iron proteins is responsible for a wide range of functions within several life forms, such as 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). When the body is fighting a bacterial infection, the body sequesters iron inside of cells (mostly stored in the storage molecule ferritin) so that it cannot be used by bacteria.

Iron distribution is heavily regulated in mammals, both as a defense against bacterial infection and because of the potential biological toxicity of iron. The iron absorbed from the duodenum binds to transferrin, and is carried by blood to different cells. There it gets incorporated, by an as yet unknown mechanism, into target proteins.[3]. A lengthier article on the system of human iron regulation can be found in the article on human iron metabolism.

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Nutrition and dietary sources

Good sources of dietary iron include meat, fish, poultry, lentils, beans, leaf vegetables, tofu, chickpeas, black-eyed peas, potatoes with skin, bread made from completely wholemeal flour, molasses and farina.

Iron provided by dietary supplements is often found as Iron (II) fumarate. Iron sulfate is as well absorbed, and less expensive. The most bioavailable form of iron supplement (ten to fifteen times more bioavailable than any other) is iron amino acid chelate. [4] The RDA for iron varies considerably based on the age, gender, and source of dietary iron (heme-based iron has higher bioavailability)[5]. Also note the section below on precautions.

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Iron in organic synthesis

The usage of iron metal filings in organic synthesis is mainly for the reduction of nitro compounds.[6] Additionally, iron has been used for desulfurizations[7], reduction of aldehydes[8], and the deoxygenation of amine oxides[9].

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Precautions

Excessive iron is toxic to humans, because excess ferrous iron reacts with peroxides in the body, producing free radicals. Iron becomes toxic when it exceeds the amount of transferrin needed to bind free iron. In excess, uncontrollable quantities of free radicals are produced.

Iron uptake is tightly regulated by the human body, which has no physiologic means of excreting iron and regulates iron solely by regulating uptake. However, too much ingested iron can damage the cells of the gastrointestinal tract directly, and may enter the bloodstream by damaging the cells that would otherwise regulate its entry. Once there, it causes damage to cells in the heart, liver and elsewhere. This can cause serious problems, including the potential of death from overdose, and long-term organ damage in survivors.

Humans experience iron toxicity above 20 milligrams of iron for every kilogram of weight, and 60 milligrams per kilogram is a lethal dose.[10] Over-consumption of iron, often the result of children eating large quantitities of ferrous sulfate tablets intended for adult consumption, is the most common toxicological cause 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.

If iron intake is excessive iron overload disorders can sometimes result, such as hemochromatosis. Iron overload disorders require a genetic inability to regulate iron uptake; however, many people have a genetic susceptibility to iron overload without realizing it and without knowing a family history of the problem. For this reason, people should not take iron supplements unless they suffer from iron deficiency and have consulted a doctor. Blood donors are at special risk of low iron levels and are often recommended to supplement their iron intake.

The medical management of iron toxicity is complex. One element of the medical approach is a specific chelating agent called deferoxamine, used to bind and expel excess iron from the body in case of iron toxicity.

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References

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Bibliography

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Cited references

  1. http://www.webelements.com/webelements/elements/text/Fe/xtal.html
  2. Dauphas, N. & Rouxel, O. 2006. Mass spectrometry and natural variations of iron isotopes. Mass Spectrometry Reviews, 25, 515-550
  3. Tracey A. Rouault. How Mammals Acquire and Distribute Iron Needed for Oxygen-Based Metabolism. Retrieved on 2006-06-19.
  4. Ashmead, H. DeWayne (1989). 'Conversations on Chelation and Mineral Nutrition'. Keats Publishing. ISBN 0-87983-501-X.
  5. Dietary Reference Intakes: Elements (PDF).
  6. Fox, B. A.; Threlfall, T. L. Organic Syntheses, Coll. Vol. 5, p.346 (1973); Vol. 44, p.34 (1964). (Article)
  7. Blomquist, A. T.; Dinguid, L. I. J. Org. Chem. 1947, 12, 718 & 723.
  8. Clarke, H. T.; Dreger, E. E. Org. Syn., Coll. Vol. 1, p.304 (1941); Vol. 6, p.52 (1926). (Article)
  9. den Hertog, J.; Overhoff, J. Recl. Trav. Chim. Pays-Bas 1950, 69, 468.
  10. Toxicity, Iron. Emidicine. Retrieved on 2006-06-19.
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See also

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External links

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