document_title: Iron

question_text: what is the purest form of iron called

question_tokens: what is the purest form of iron called

text: Iron

text: Characteristics

text: Mechanical properties

text:
Characteristic values of tensile strength (TS) and Brinell hardness (BH) of different forms of iron.[6][7]

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

Carbon-containing iron
140
40

Pure, single-crystal iron
10
3

text: The mechanical properties of iron and its alloys can be evaluated using a variety of tests, including the Brinell test, Rockwell test and the Vickers hardness test. The data on iron is so consistent that it is often used to calibrate measurements or to compare tests.[7][8] However, the mechanical properties of iron are significantly affected by the sample's purity: pure, single crystals of iron are actually softer than aluminium,[6] and the purest industrially produced iron (99.99%) has a hardness of 20–30 Brinell.[9] An increase in the carbon content will cause a significant increase in the hardness and tensile strength of iron. Maximum hardness of 65 Rc is achieved with a 0.6% carbon content, although the alloy has low tensile strength.[10] Because of the softness of iron, it is much easier to work with than its heavier congeners ruthenium and osmium.[11]

text: Because of its significance for planetary cores, the physical properties of iron at high pressures and temperatures have also been studied extensively. The form of iron that is stable under standard conditions can be subjected to pressures up to ca. 15 GPa before transforming into a high-pressure form, as described in the next section.

text: Phase diagram and allotropes

text: Iron represents an example of allotropy in a metal. There are at least four allotropic forms of iron, known as α, γ, δ, and ε; at very high pressures and temperatures, some controversial experimental evidence exists for a stable β phase.[12]

text: As molten iron cools past its freezing point of 1538 °C, it crystallizes into its δ allotrope, which has a body-centered cubic (bcc) crystal structure. As it cools further to 1394 °C, it changes to its γ-iron allotrope, a face-centered cubic (fcc) crystal structure, or austenite. At 912 °C and below, the crystal structure again becomes the bcc α-iron allotrope, or ferrite. Finally, at 770 °C (the Curie point, Tc) iron's magnetic ordering changes from paramagnetic to ferromagnetic. 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 have the same axis orientation; however, the electrons of neighboring domains have other orientations with the result of mutual cancellation and no magnetic field. In magnetized iron, the electronic spins of the domains are aligned and the magnetic effects are reinforced. Although each domain contains billions of atoms, they are very small, about 10 micrometres across.[13] This happens because the two unpaired electrons on each iron atom are in the dz2 and dx2 − y2 orbitals, which do not point directly at the nearest neighbors in the body-centered cubic lattice and therefore do not participate in metallic bonding; thus, they can interact magnetically with each other so that their spins align.[14]

text: At pressures above approximately 10 GPa and temperatures of a few hundred kelvin or less, α-iron changes into a hexagonal close-packed (hcp) structure, which is also known as ε-iron; the higher-temperature γ-phase also changes into ε-iron, but does so at higher pressure. The β-phase, if it exists, would appear at pressures of at least 50 GPa and temperatures of at least 1500 K and have an orthorhombic or a double hcp structure.[12] These high-pressure phases of iron are important as endmember models for the solid parts of planetary cores. The inner core of the Earth is generally presumed to be an iron-nickel alloy with ε (or β) structure.[15] Somewhat confusingly, the term "β-iron" is sometimes also used to refer to α-iron above its Curie point, when it changes from being ferromagnetic to paramagnetic, even though its crystal structure has not changed.[14]

text: The melting point of iron is experimentally well defined for pressures less than 50 GPa. For greater pressures, studies put the γ-ε-liquid triple point at pressures that differ by tens of gigapascals and 1000 K in the melting point. Generally speaking, molecular dynamics computer simulations of iron melting and shock wave experiments suggest higher melting points and a much steeper slope of the melting curve than static experiments carried out in diamond anvil cells.[16] The melting and boiling points of iron, along with its enthalpy of atomization, are lower than those of the earlier 3d elements from scandium to chromium, showing the lessened contribution of the 3d electrons to metallic bonding as they are attracted more and more into the inert core by the nucleus;[17] however, they are higher than the values for the previous element manganese because that element has a half-filled 3d subshell and consequently its d-electrons are not easily delocalized. This same trend appears for ruthenium but not osmium.[11]

text: Isotopes

text: 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. Of these stable isotopes, only 57Fe has a nuclear spin (− 1⁄2). The nuclide 54Fe theoretically can undergo double electron capture to 54Cr, but the process has never been observed and only a lower limit on the half-life of 3.1×1022 years has been established.[18]

text: 60Fe is an extinct radionuclide of long half-life (2.6 million years).[19] It is not found on Earth, but its ultimate decay product is its granddaughter, the stable nuclide 60Ni.[18] Much of the past work on isotopic composition of iron has focused on the nucleosynthesis of 60Fe through studies of meteorites and ore formation. In the last decade, advances in mass spectrometry have allowed the detection and quantification of minute, naturally occurring variations in the ratios of the stable isotopes of iron. Much of this work is driven by the Earth and planetary science communities, although applications to biological and industrial systems are emerging.[20]

text: In phases of the meteorites Semarkona and Chervony Kut, a correlation between the concentration of 60Ni, the granddaughter of 60Fe, and the abundance of the stable iron isotopes provided evidence for the existence of 60Fe at the time of formation of the Solar System. Possibly the energy released by the decay of 60Fe, along with that released by 26Al, contributed to the remelting and differentiation of asteroids after their formation 4.6 billion years ago. The abundance of 60Ni present in extraterrestrial material may bring further insight into the origin and early history of the Solar System.[21]

text: The most abundant iron isotope 56Fe is of particular interest to nuclear scientists because it represents the most common endpoint of nucleosynthesis.[22] Since 56Ni (14 alpha particles) is easily produced from lighter nuclei in the alpha process in nuclear reactions in supernovae (see silicon burning process), it is the endpoint of fusion chains inside extremely massive stars, since addition of another alpha particle, resulting in 60Zn, requires a great deal more energy. This 56Ni, which has a half-life of about 6 days, is created in quantity in these stars, but soon decays by two successive positron emissions within supernova decay products in the supernova remnant gas cloud, first to radioactive 56Co, and then to stable 56Fe. As such, 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.[23] It is also very common in the universe, relative to other stable metals of approximately the same atomic weight.[23][24] Iron is the sixth most abundant element in the Universe, and the most common refractory element.[25]

text: Although a further tiny energy gain could be extracted by synthesizing 62Ni, which has a marginally higher binding energy than 56Fe, conditions in stars are unsuitable for this process. Element production in supernovas and distribution on Earth greatly favor iron over nickel, and in any case, 56Fe still has a lower mass per nucleon than 62Ni due to its higher fraction of lighter protons.[26] Hence, elements heavier than iron require a supernova for their formation, involving rapid neutron capture by starting 56Fe nuclei.[23]

text: In the far future of the universe, assuming that proton decay does not occur, cold fusion occurring via quantum tunnelling would cause the light nuclei in ordinary matter to fuse into 56Fe nuclei. Fission and alpha-particle emission would then make heavy nuclei decay into iron, converting all stellar-mass objects to cold spheres of pure iron.[27]

text: Occurrence

text: Metallic or native 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, after oxygen, silicon, and aluminium.[28][29] Most of the iron in the crust is found combined with oxygen as iron oxide minerals such as hematite (Fe2O3), magnetite (Fe3O4), and siderite (FeCO3). Many igneous rocks also contain the sulfide minerals pyrrhotite and pentlandite.[30][31]

text: Ferropericlase (Mg,Fe)O, a solid solution of periclase (MgO) and wüstite (FeO), makes up about 20% of the volume of the lower mantle of the Earth, which makes it the second most abundant mineral phase in that region after silicate perovskite (Mg,Fe)SiO3; it also is the major host for iron in the lower mantle.[32] At the bottom of the transition zone of the mantle, the reaction γ-(Mg,Fe)2[SiO4] ↔ (Mg,Fe)[SiO3] + (Mg,Fe)O transforms γ-olivine into a mixture of perovskite and ferropericlase and vice versa. In the literature, this mineral phase of the lower mantle is also often called magnesiowüstite.[33] Silicate perovskite may form up to 93% of the lower mantle,[34] and the magnesium iron form, (Mg,Fe)SiO3, is considered to be the most abundant mineral in the Earth, making up 38% of its volume.[35]

text: 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 alternating with bands of iron-poor shale and chert. The banded iron formations were laid down in the time between 3,700 million years ago and 1,800 million years ago.[36][37]

text: The mentioned iron compounds have been used as pigments (compare ochre) since historical time and contribute as well to the color of various geological formations, e.g. the Bundsandstein (British Bunter, colored sandstein).[38] In the case of the Eisensandstein (a jurassic 'iron sandstone', e.g. from Donzdorf) in Germany[39] and Bath stone in the UK, iron pigments contribute to the yellowish color of large amounts of historical buildings and sculptures.[40] The proverbial red color of the surface of Mars is derived from an iron oxide-rich regolith.[41]

text: Significant amounts of iron occur in the iron sulfide mineral pyrite (FeS2), but it is difficult to extract iron from it and it is therefore not used. In fact, iron is so common that production generally focuses only on ores with very high quantities of it. During weathering, iron tends to leach from sulfide deposits as the sulfate and from silicate deposits as the bicarbonate. Both of these are oxidized in aqueous solution and precipitate in even mildly elevated pH as iron(III) oxide.[42]

text: 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.[43] According to the International Resource Panel's Metal Stocks in Society report, the global stock of iron in use in society is 2200 kg per capita. Much of this is in more-developed countries (7000–14000 kg per capita) rather than less-developed countries (2000 kg per capita).[44]

text: Chemistry and compounds

text:
Oxidation
state
Representative compound

−2 (d10)
Disodium tetracarbonylferrate (Collman's reagent)

−1 (d9)
Fe
2(CO)2−
8

0 (d8)
Iron pentacarbonyl

1 (d7)
Cyclopentadienyliron dicarbonyl dimer ("Fp2")

2 (d6)
Ferrous sulfate, ferrocene

3 (d5)
Ferric chloride, ferrocenium tetrafluoroborate

4 (d4)
Fe(diars)
2Cl2+
2

5 (d3)
FeO3−
4

6 (d2)
Potassium ferrate

7 (d1)
[FeO4]– (matrix isolation, 4K)

text: Iron shows the characteristic chemical properties of the transition metals, namely the ability to form variable oxidation states differing by steps of one and a very large coordination and organometallic chemistry: indeed, it was the discovery of an iron compound, ferrocene, that revolutionalized the latter field in the 1950s.[45] Iron is sometimes considered as a prototype for the entire block of transition metals, due to its abundance and the immense role it has played in the technological progress of humanity.[46] Its 26 electrons are arranged in the configuration [Ar]3d64s2, of which the 3d and 4s electrons are relatively close in energy, and thus it can lose a variable number of electrons and there is no clear point where further ionization becomes unprofitable.[11]

text: 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, although this is very easily reduced. Although iron(VIII) oxide (FeO4) has been claimed, the report could not be reproduced and such a species (at least with iron in its +8 oxidation state) has been found to be improbable computationally.[47] However, one form of anionic [FeO4]– with iron in its +7 oxidation state, along with an iron(V)-peroxo isomer, has been detected by infrared spectroscopy at 4K after cocondensation of laser-ablated Fe atoms with a mixture of O2/Ar.[48] Iron(IV) is a common intermediate in many biochemical oxidation reactions.[49][50] 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.[51] 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).[50] The latter is used as the traditional "blue" in blueprints.[52]

text: Iron is the first of the transition metals that cannot reach its group oxidation state of +8, although its heavier congeners ruthenium and osmium can, with ruthenium having more difficulty than osmium.[14] Ruthenium exhibits an aqueous cationic chemistry in its low oxidation states similar to that of iron, but osmium does not, favoring high oxidation states in which it forms anionic complexes.[14] In fact, in this second half of the 3d transition series, vertical similarities down the groups compete with the horizontal similarities of iron with its neighbors cobalt and nickel in the periodic table, which are also ferromagnetic at room temperature and share similar chemistry. As such, iron, cobalt, and nickel are sometimes grouped together as the iron triad.[46]

text: 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.[50]

text: 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.[53]

text: Iron is by far the most reactive element in its group; it is pyrophoric when finely divided and dissolves easily in dilute acids, giving Fe2+. However, it does not react with concentrated nitric acid and other oxidizing acids due to the formation of an impervious oxide layer, which can nevertheless react with hydrochloric acid.[14]

text: Binary compounds

text:
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. Despite their names, they are actually all non-stoichiometric compounds whose compositions may vary.[54] 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.[50] It is not an iron(IV) compound, but is actually an iron(II) polysulfide containing Fe2+ and S2−
2 ions in a distorted sodium chloride structure.[54]

text: 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 hydrohalic acid to give the corresponding hydrated salts.[50]

text: Iron reacts with fluorine, chlorine, and bromine to give the corresponding ferric halides, ferric chloride being the most common.[55]

text: Ferric iodide is an exception, being thermodynamically unstable due to the oxidizing power of Fe3+ and the high reducing power of I−:[55]

text: Nevertheless, milligram amounts of ferric iodide, a black solid, may still be prepared through the reaction of iron pentacarbonyl with iodine and carbon monoxide in the presence of hexane and light at the temperature of −20 °C, making sure that the system is well sealed off from air and water.[55]

text: Solution chemistry

text: The standard reduction potentials in acidic aqueous solution for some common iron ions are given below:[14]

text:
Fe2+ + 2 e−
⇌ Fe
E0 = −0.447 V

Fe3+ + 3 e−
⇌ Fe
E0 = −0.037 V

FeO2−
4 + 8 H+ + 3 e−
⇌ Fe3+ + 4 H2O
E0 = +2.20 V

text: The red-purple tetrahedral ferrate(VI) anion is such a strong oxidizing agent that it oxidizes nitrogen and ammonia at room temperature, and even water itself in acidic or neutral solutions:[55]

text: The Fe3+ ion has a large simple cationic chemistry, although the pale-violet hexaquo ion [Fe(H2O)6]3+ is very readily hydrolyzed when pH increases above 0 as follows:[56]

text:
[Fe(H2O)6]3+
⇌ [Fe(H2O)5(OH)]2+ + H+
K = 10−3.05 mol dm−3

[Fe(H2O)5(OH)]2+
⇌ [Fe(H2O)4(OH)2]+ + H+
K = 10−3.26 mol dm−3

2 [Fe(H2O)6]3+
⇌ [Fe(H
2O)
4(OH)]4+
2 + 2 H+ + 2 H2O
K = 10−2.91 mol dm−3

text: As pH rises above 0 the above yellow hydrolyzed species form and as it rises above 2–3, reddish-brown hydrous iron(III) oxide precipitates out of solution. Although Fe3+ has an d5 configuration, its absorption spectrum is not like that of Mn2+ with its weak, spin-forbidden d–d bands, because Fe3+ has higher positive charge and is more polarizing, lowering the energy of its ligand-to-metal charge transfer absorptions. Thus, all the above complexes are rather strongly colored, with the single exception of the hexaquo ion – and even that has a spectrum dominated by charge transfer in the near ultraviolet region.[56] On the other hand, the pale green iron(II) hexaquo ion [Fe(H2O)6]2+ does not undergo appreciable hydrolysis. Carbon dioxide is not evolved when carbonate anions are added, which instead results in white iron(II) carbonate being precipitated out. In excess carbon dioxide this forms the slightly soluble bicarbonate, which occurs commonly in groundwater, but it oxidises quickly in air to form iron(III) oxide that accounts for the brown deposits present in a sizeable number of streams.[57]

text: Coordination compounds

text: Many coordination compounds of iron are known. A typical six-coordinate anion is hexachloroferrate(III), [FeCl6]3−, found in the mixed salt tetrakis(methylammonium) hexachloroferrate(III) chloride.[58][59] Complexes with multiple bidentate ligands have geometric isomers. For example, the trans-chlorohydridobis(bis-1,2-(diphenylphosphino)ethane)iron(II) complex is used as a starting material for compounds with the Fe(dppe)2 moiety.[60][61] The ferrioxalate ion with three oxalate ligands (shown at right) displays helical chirality with its two non-superposable geometries labelled Λ (lambda) for the left-handed screw axis and Δ (delta) for the right-handed screw axis, in line with IUPAC conventions.[56] Potassium ferrioxalate is used in chemical actinometry and along with its sodium salt undergoes photoreduction applied in old-style photographic processes. The dihydrate of iron(II) oxalate has a polymeric structure with co-planar oxalate ions bridging between iron centres with the water of crystallisation located forming the caps of each octahedron, as illustrated below.[62]

text: Prussian blue, Fe4[Fe(CN)6]3, is the most famous of the cyanide complexes of iron. Its formation can be used as a simple wet chemistry test to distinguish between aqueous solutions of Fe2+ and Fe3+ as they react (respectively) with potassium ferricyanide and potassium ferrocyanide to form Prussian blue.[50]

text: Iron(III) complexes are quite similar to those of chromium(III) with the exception of iron(III)'s preference for O-donor instead of N-donor ligands. The latter tend to be rather more unstable than iron(II) complexes and often dissociate in water. Many Fe–O complexes show intense colors and are used as tests for phenols or enols. For example, in the ferric chloride test, used to determine the presence of phenols, iron(III) chloride reacts with a phenol to form a deep violet complex:[56]

text: Among the halide and pseudohalide complexes, fluoro complexes of iron(III) are the most stable, with the colorless [FeF5(H2O)]2− being the most stable in aqueous solution. Chloro complexes are less stable and favor tetrahedral coordination as in [FeCl4]−; finally, [FeBr4]− and [FeI4]− reduce themselves easily to iron(II). Thiocyanate is a common test for the presence of iron(III) as it forms the blood-red [Fe(SCN)(H2O)5]2+. Like manganese(II), most iron(III) complexes are high-spin, the exceptions being those with ligands that are high in the spectrochemical series such as cyanide. An example of a low-spin iron(III) complex is [Fe(CN)6]3−. The cyanide ligands may easily be detached in [Fe(CN)6]3−, and hence this complex is poisonous, unlike the iron(II) complex [Fe(CN)6]4− found in Prussian blue,[56] which does not release hydrogen cyanide except when dilute acids are added.[57] Iron shows a great variety of electronic spin states, including every possible spin quantum number value for a d-block element from 0 (diamagnetic) to  5⁄2 (5 unpaired electrons). This value is always half the number of unpaired electrons. Complexes with zero to two unpaired electrons are considered low-spin and those with four or five are considered high-spin.[54]

text: Iron(II) complexes are less stable than iron(III) complexes but the preference for O-donor ligands is less marked, so that for example [Fe(NH3)6]2+ is known while [Fe(NH3)6]3+ is not. They have a tendency to be oxidized to iron(III) but this can be moderated by low pH and the specific ligands used.[57]

text: Organometallic compounds

text: Cyanide complexes are technically organometallic but more important are carbonyl complexes and sandwich and half-sandwich compounds. 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.[63]

text: Ferrocene was an extremely important compound in the early history of the branch of organometallic chemistry, and to this day iron is still one of the most important metals in this field.[64] It was first synthesised in 1951 during an attempt to prepare the fulvalene (C10H8) by oxidative dimerization of cyclopentadiene; the resultant product was found to have molecular formula C10H10Fe and reported to exhibit "remarkable stability".[65] The discovery sparked substantial interest in the field of organometallic chemistry,[66][67] in part because the structure proposed by Pauson and Kealy (shown at right) was inconsistent with then-existing bonding models and did not explain its unexpected stability. Consequently, the initial challenge was to definitively determine the structure of ferrocene in the hope that its bonding and properties would then be understood. The shockingly novel sandwich structure, [Fe(η5-C5H5)2],[66] was deduced and reported independently by three groups in 1952: Robert Burns Woodward and Geoffrey Wilkinson investigated the reactivity in order to determine the structure[68] and demonstrated that ferrocene undergoes similar reactions to a typical aromatic molecule (such as benzene),[69] Ernst Otto Fischer deduced the sandwich structure and also began synthesising other metallocenes including cobaltocene;[70] Eiland and Pepinsky provided X-ray crystallographic confirmation of the sandwich structure.[71]

text: Applying valence bond theory to ferrocene by considering an Fe2+ centre and two cyclopentadienide anions (C5H5−), which are known to be aromatic according to Hückel's rule and hence highly stable, allowed correct prediction of the geometry of the molecule. Once molecular orbital theory was successfully applied and the Dewar-Chatt-Duncanson model proposed,[72] the reasons for ferrocene's remarkable stability became clear.[73] Ferrocene was not the first organometallic compound known – Zeise's salt, K[PtCl3(C2H4)]·H2O was reported in 1831[74][75] and Mond's discovery of Ni(CO)4 occurred in 1888,[76] but it was ferrocene's discovery that began organometallic chemistry as a separate area of chemistry. It was so important that Wilkinson and Fischer shared the 1973 Nobel Prize for Chemistry "for their pioneering work, performed independently, on the chemistry of the organometallic, so called sandwich compounds".[77] Ferrocene itself can be used as the backbone of a ligand, e.g. 1,1'-bis(diphenylphosphino)ferrocene (dppf). Ferrocene can itself be oxidized to the ferrocenium cation (Fc+); the ferrocene/ferrocenium couple is often used as a reference in electrochemistry.[67]

text: Metallocenes like ferrocene can be prepared by reaction of freshly-cracked cyclopentadiene with iron(II) chloride and a weak base.[78] It is an aromatic substance and undergoes substitution reactions rather than addition reactions on the cyclopentadienyl ligands. For example, Friedel-Crafts acylation of ferrocene with acetic anhydride yields acetylferrocene[79] just as acylation of benzene yields acetophenone under similar conditions.

text: Iron-centered organometallic species are used as catalysts. The Knölker complex, for example, is a transfer hydrogenation catalyst for ketones.[80]

text: Etymology

text: As iron has been in use for such a long time, it has many different names in different languages. The source of its chemical symbol Fe is the Latin word ferrum, and its descendants are the names of the element in the Romance languages (for example, French fer, Spanish hierro, and Italian and Portuguese ferro).[81] The word ferrum itself possibly comes from the Semitic languages, via Etruscan, from a root that also gave rise to Old English bræs "brass".[82] The English word iron derives ultimately from Proto-Germanic *isarnan, which is also the source of the German name Eisen. It was most likely borrowed from Celtic *isarnon, which ultimately comes from Proto-Indo-European *is-(e)ro- "powerful, holy" and finally *eis "strong", referencing iron's strength as a metal.[83] Kluge relates *isarnon to Illyric and Latin ira, 'wrath'[4]) The Balto-Slavic names for iron (for example, Russian железо [zhelezo]) are the only ones to come directly from the Proto-Indo-European *ghelgh- "iron".[84] In many of these languages, the word for iron may also be used to denote other objects made of iron or steel, or figuratively because of the hardness and strength of the metal.[85] The Chinese tiě (traditional 鐵; simplified 铁) derives from Proto-Sino-Tibetan *hliek,[86] and was borrowed into Japanese as 鉄 tetsu, which also has the native reading kurogane "black metal" (similar to how iron is referenced in the English word blacksmith).[87]

text: History

text: Wrought iron

text: Iron belongs to the elements undoubtedly known to the ancient world.[88] It has been worked, or wrought, for millennia. However, iron objects of great age are much rarer than objects made of gold or silver due to the ease with which iron corrodes.[89]

text: Beads made from meteoric iron in 3500 BC or earlier were found in Gerzah, Egypt by G. A. Wainwright.[90] The beads contain 7.5% nickel, which is a signature of meteoric origin since iron found in the Earth's crust generally has only minuscule nickel impurities. Meteoric iron was highly regarded due to its origin in the heavens and was often used to forge weapons and tools.[90] For example, a dagger made of meteoric iron was found in the tomb of Tutankhamun, containing similar proportions of iron, cobalt, and nickel to a meteorite discovered in the area, deposited by an ancient meteor shower.[91][92][93] Items that were likely made of iron by Egyptians date from 3000 to 2500 BC.[89] Meteoritic iron is comparably soft and ductile and easily forged by cold working but may get brittle when heated because of the nickel content.[94]

text: 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 3000 and 2700 BC.[95] The Hittites established an empire in north-central Anatolia around 1600 BC. They appear to be the first to understand the production of iron from its ores and regard it highly in their society.[96] The Hittites 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.[95] The subsequent period is called the Iron Age.

text: Artifacts of smelted iron are found in India dating from 1800 to 1200 BC,[97] and in the Levant from about 1500 BC (suggesting smelting in Anatolia or the Caucasus).[98][99] Alleged references (compare history of metallurgy in South Asia) to iron in the Indian Vedas have been used for claims of a very early usage of iron in India respectively to date the texts as such. The rigveda term ayas (metal) probably refers to copper and bronze, while iron or śyāma ayas, literally "black metal", first is mentioned in the post-rigvedic Atharvaveda.[100]

text: There is some archaeological evidence of iron being smelted in Zimbabwe and southeast Africa as early as the eighth century BC.[101] Iron working was introduced to Greece in the late 11th century BC, from which it spread quickly throughout Europe.[102]

text: The spread of ironworking in Central and Western Europe is associated with Celtic expansion. According to Pliny the Elder, iron use was common in the Roman era.[90] The annual iron output of the Roman Empire is estimated at 84750 t,[103] while the similarly populous and contemporary Han China produced around 5000 t.[104] In China, iron only appears circa 700–500 BC.[105] Iron smelting may have been introduced into China through Central Asia.[106] The earliest evidence of the use of a blast furnace in China dates to the 1st century AD,[107] and cupola furnaces were used as early as the Warring States period (403–221 BC).[108] Usage of the blast and cupola furnace remained widespread during the Song and Tang Dynasties.[109]

text: During the Industrial Revolution in Britain, Henry Cort began refining iron from pig iron to wrought iron (or bar iron) using innovative production systems. In 1783 he patented the puddling process for refining iron ore. It was later improved by others, including Joseph Hall.[110]

text: Cast iron

text: Cast iron was first produced in China during 5th century BC,[111] but was hardly in Europe until the medieval period.[112][113] The earliest cast iron artifacts were discovered by archaeologists in what is now modern Luhe County, Jiangsu in China. Cast iron was used in ancient China for warfare, agriculture, and architecture.[114] 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.[115]

text: Medieval blast furnaces were about 10 feet (3.0 m) tall and made of fireproof brick; forced air was usually provided by hand-operated bellows.[113] Modern blast furnaces have grown much bigger, with hearths fourteen meters in diameter that allow them to produce thousands of tons of iron each day, but essentially operate in much the same way as they did during medieval times.[115]

text: In 1709, Abraham Darby I established a coke-fired blast furnace to produce cast iron, replacing charcoal, although continuing to use blast furnaces. The ensuing availability of inexpensive iron was one of the factors leading to the Industrial Revolution. Toward the end of the 18th century, cast iron began to replace wrought iron for certain purposes, because it was cheaper. Carbon content in iron was not implicated as the reason for the differences in properties of wrought iron, cast iron, and steel until the 18th century.[95]

text: Since iron was becoming cheaper and more plentiful, it also became a major structural material following the building of the innovative first iron bridge in 1778. This bridge still stands today as a monument to the role iron played in the Industrial Revolution. Following this, iron was used in rails, boats, ships, aqueducts, and buildings, as well as in iron cylinders in steam engines.[115] Railways have been central to the formation of modernity and ideas of progress[116] and various languages (e.g. French, Spanish, Italien and German) refer to railways as iron road.

text: Steel

text: 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 Persia were making good steel by 1000 BC.[95] Then improved versions, Wootz steel by India and Damascus steel were developed around 300 BC and AD 500 respectively. These methods were specialized, and so steel did not become a major commodity until the 1850s.[117]

text: New methods of producing it by carburizing bars of iron in the cementation process were devised in the 17th century. 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 in large quantities.[118]

text: Foundations of modern chemistry

text: In 1774, Antoine Lavoisier used the reaction of water steam with metallic iron inside an incandescent iron tube to produce hydrogen in his experiments leading to the demonstration of the conservation of mass, which was instrumental in changing chemistry from a qualitative science to a quantitative one.[119]

text: Symbolic role

text: Iron plays a certain role in mythology and has found various usage as a metaphor and in folklore. The Greek poet Hesiod's Works and Days (lines 109–201) lists different ages of man named after metals like gold, silver, bronze and iron to account for successive ages of humanity.[120] The iron age was closely related with Rome, and in Ovid's Metamorphoses

text: The Virtues, in despair, quit the earth; and the depravity of man becomes universal and complete. Hard steel succeeded then.

text: An example of the importance of iron's symbolic role may be found in the German Campaign of 1813. Frederick William III commissioned then the first Iron Cross as military decoration. Berlin iron jewellery reached its peak production between 1813 and 1815, when the Prussian royal family urged citizens to donate gold and silver jewellery for military funding. The inscription Gold gab ich für Eisen (I gave gold for iron) was used as well in later war efforts.[121]

text: Production of metallic iron

text: Industrial routes

text: The production of iron or steel is a process consisting of two main stages. In the first stage pig iron is produced in a blast furnace. Alternatively, it may be directly reduced. In the second stage, pig iron is converted to wrought iron, steel, or cast iron.[122]

text: For a few limited purposes when it is needed, pure iron is produced in the laboratory in small quantities by reducing the pure oxide or hydroxide with hydrogen, or forming iron pentacarbonyl and heating it to 250 °C so that it decomposes to form pure iron powder.[42] Another method is electrolysis of ferrous chloride onto an iron cathode.[123]

text: Blast furnace processing

text: Industrial iron production starts with iron ores, principally hematite, which has a nominal formula Fe2O3, and magnetite, with the formula Fe3O4. These ores are reduced to the metal in a carbothermic reaction, i.e. by treatment with carbon. The conversion is typically conducted in a blast furnace at temperatures of about 2000 °C. Carbon is provided in the form of coke. The process also contains a flux such as limestone, which is used to remove silicaceous minerals in the ore, which would otherwise clog the furnace. The coke and limestone are fed into the top of the furnace, while a massive blast of air heated to 900 °C, about 4 tons per ton of iron,[113] is forced into the furnace at the bottom.[122]

text: In the furnace, the coke reacts with oxygen in the air blast to produce carbon monoxide:[122]

text: The carbon monoxide reduces the iron ore (in the chemical equation below, hematite) to molten iron, becoming carbon dioxide in the process:[122]

text: Some iron in the high-temperature lower region of the furnace reacts directly with the coke:[122]

text: The flux present to melt impurities in the ore is principally limestone (calcium carbonate) and dolomite (calcium-magnesium carbonate). Other specialized fluxes are used depending on the details of the ore. In the heat of the furnace the limestone flux decomposes to calcium oxide (also known as quicklime):[122]

text: Then calcium oxide combines with silicon dioxide to form a liquid slag.[122]

text: 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.[113]

text: Direct iron reduction

text: Owing to environmental concerns, alternative methods of processing iron have been developed. "Direct iron reduction" reduces iron ore to a ferrous lump called "sponge" iron or "direct" iron that is suitable for steelmaking.[113] Two main reactions comprise the direct reduction process:

text: Natural gas is partially oxidized (with heat and a catalyst):[113]

text: These gases are then treated with iron ore in a furnace, producing solid sponge iron:[113]

text: Silica is removed by adding a limestone flux as described above.[113]

text: Thermite

text: Iron is a byproduct of burning a mixture of aluminum powder and rust powder.

text: Further processes

text: 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.[122] Removing the other impurities results in cast iron, which is used to cast articles in foundries such as stoves, pipes, radiators, lamp-posts and rails.[122]

text: 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.[115]

text: 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.[124]

text: Applications

text: Metallurgical

text: Iron is the most widely used of all the metals, accounting for over 90% 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 combined with alloying elements to make steel.[125]

text: Ferrite (α-iron) is a fairly soft metal that can dissolve only a small concentration of carbon (no more than 0.021% by mass at 910 °C).[126] Austenite (γ-iron) is similarly soft and metallic but can dissolve considerably more carbon (as much as 2.04% by mass at 1146 °C). This form of iron is used in the type of stainless steel used for making cutlery, and hospital and food-service equipment.[13]

text: Commercially available iron is classified based on purity and the abundance of additives. Pig iron has 3.5–4.5% carbon[127] 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. The reduction of contaminants in pig iron that negatively affect material properties, such as sulfur and phosphorus, yields cast iron containing 2–4% carbon, 1–6% silicon, and small amounts of manganese.[122] Pig iron 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.[14] Its mechanical properties vary greatly and depend on the form the carbon takes in the alloy.[11]

text: "White" cast irons contain their carbon in the form of cementite, or iron carbide (Fe3C).[11] This hard, brittle compound dominates the mechanical properties of white cast irons, rendering them hard, but unresistant to shock. The broken surface of a white cast iron is full of fine facets of the broken iron-carbide, a very pale, silvery, shiny material, hence the appellation. Cooling a mixture of iron with 0.8% carbon slowly below 723 °C to room temperature results in separate, alternating layers of cementite and ferrite, which is soft and malleable and is called pearlite for its appearance. Rapid cooling, on the other hand, does not allow time for this separation and creates hard and brittle martensite. The steel can then be tempered by reheating to a temperature in between, changing the proportions of pearlite and martensite. The end product below 0.8% carbon content is a pearlite-ferrite mixture, and that above 0.8% carbon content is a pearlite-cementite mixture.[11]

text: In gray iron the carbon exists as separate, fine flakes of graphite, and also renders the material brittle due to the sharp edged flakes of graphite that produce stress concentration sites within the material.[128] 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, reducing the stress concentrations and vastly increasing the toughness and strength of the material.[128]

text: Wrought iron contains less than 0.25% carbon but large amounts of slag that give it a fibrous characteristic.[127] 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 within 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.

text:
Iron production 2009 (million tonnes)[129]

Country
Iron ore
Pig iron
Direct iron
Steel

China
1,114.9
549.4

573.6

Australia
393.9
4.4

5.2

Brazil
305.0
25.1
0.011
26.5

Japan

66.9

87.5

India
257.4
38.2
23.4
63.5

Russia
92.1
43.9
4.7
60.0

Ukraine
65.8
25.7

29.9

South Korea
0.1
27.3

48.6

Germany
0.4
20.1
0.38
32.7

World
1,594.9
914.0
64.5
1,232.4

text: Mild steel corrodes more readily than wrought iron, but is cheaper and more widely available. Carbon steel contains 2.0% carbon or less,[130] 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.[130][131][132]

text: 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.[133]

text: 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, a cost amounting to over 1% of the world's economy.[134] Painting, galvanization, passivation, plastic coating and bluing are all used to protect iron from rust by excluding water and oxygen or by cathodic protection. The mechanism of the rusting of iron is as follows:[134]

text: The electrolyte is usually iron(II) sulfate in urban areas (formed when atmospheric sulfur dioxide attacks iron), and salt particles in the atmosphere in seaside areas.[134]

text: Iron compounds

text: Although the dominant use of iron is in metallurgy, iron compounds are also pervasive in industry. 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.[135] Powdered iron in an acidic solvent was used in the Bechamp reduction the reduction of nitrobenzene to aniline.[136]

text: 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.[137] It can also be dissolved in alcohol to form tincture of iron, which is used as a medicine to stop bleeding in canaries.[138]

text: 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. 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.[137]

text: Biological and pathological role

text: Iron is involved in numerous biological processes.[139][140] It is the most important transition metal in all living organisms.[141] Iron-proteins are found in all living organisms: archaeans, bacteria and eukaryotes, including humans. For example, the color of blood is due to hemoglobin, an iron-containing protein. As illustrated by hemoglobin, iron is often bound to cofactors, such as hemes, which are non-protein compounds, often involving metal ions, that are required for a protein's biological activity to happen. The iron-sulfur clusters are pervasive and include nitrogenase, the enzymes responsible for biological nitrogen fixation. The main roles of iron-containing proteins are the transport and storage of oxygen, as well as the transfer of electrons.[141]

text: 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 (see high-valent iron), and catalase.[142] The average adult human contains about 0.005% body weight of iron, or about four grams, of which three quarters is in hemoglobin – a level that remains constant despite only about one milligram of iron being absorbed each day,[141] because the human body recycles its hemoglobin for the iron content.[143]

text: Biochemistry

text: Iron acquisition poses a problem for aerobic organisms because ferric iron is poorly soluble near neutral pH. Thus, these organisms have developed means to absorb iron as complexes, sometimes taking up ferrous iron before oxidising it back to ferric iron. In particular, bacteria have evolved very high-affinity sequestering agents called siderophores.[144][145][146]

text: After uptake in human cells, iron storage is carefully regulated; iron ions are never "free". This is because free iron ions have a high potential for biological toxicity.[147] A major component of this regulation is the protein transferrin, which binds iron ions absorbed from the duodenum and carries it in the blood to cells.[148] Transferrin contains Fe3+ in the middle of a distorted octahedron, bonded to one nitrogen, three oxygens and a chelating carbonate anion that traps the Fe3+ ion: it has such a high stability constant that it is very effective at taking up Fe3+ ions even from the most stable complexes. At the bone marrow, transferrin is reduced from Fe3+ and Fe2+ and stored as ferritin to be incorporated into hemoglobin.[141]

text: The most commonly known and studied bioinorganic iron compounds (biological iron molecules) are the heme proteins: examples are hemoglobin, myoglobin, and cytochrome P450. These compounds participate in transporting gases, building enzymes, and transferring electrons.[141] Metalloproteins are a group of proteins with metal ion cofactors. Some examples of iron metalloproteins are ferritin and rubredoxin.[141] Many enzymes vital to life contain iron, such as catalase,[149] lipoxygenases,[150] and IRE-BP.[151]

text: Hemoglobin is an oxygen carrier that occurs in red blood cells and contributes their color, transporting oxygen in the arteries from the lungs to the muscles where it is transferred to myoglobin, which stores it until it is needed for the metabolic oxidation of glucose, which generates energy. Here the hemoglobin binds to carbon dioxide, produced when glucose is oxidized, which is transported through the veins by hemoglobin (predominantly as bicarbonate anions) back to the lungs where it is exhaled.[141] In hemoglobin, the iron is in one of four heme groups and has six possible coordination sites; four are occupied by nitrogen atoms in a porphyrin ring, the fifth by an imidazole nitrogen in a histidine residue of one of the protein chains attached to the heme group, and the sixth is reserved for the oxygen molecule it can reversibly bind to.[141] When hemoglobin is not attached to oxygen (and is then called deoxyhemoglobin), the Fe2+ ion at the center of the heme group (in the hydrophobic protein interior) is in a high-spin configuration. It is thus too large to fit inside the porphyrin ring, which bends instead into a dome with the Fe2+ ion about 55 picometers above it. In this configuration, the sixth coordination site reserved for the oxygen is blocked by another histidine residue.[141]

text: When deoxyhemoglobin picks up an oxygen molecule, this histidine residue moves away and returns once the oxygen is securely attached to form a hydrogen bond with it. This results in the Fe2+ ion switching to a low-spin configuration, resulting in a 20% decrease in ionic radius so that now it can fit into the porphyrin ring, which becomes planar.[141] (Additionally, this hydrogen bonding results in the tilting of the oxygen molecule, resulting in a Fe–O–O bond angle of around 120° that avoids the formation of Fe–O–Fe or Fe–O2–Fe bridges that would lead to electron transfer, the oxidation of Fe2+ to Fe3+, and the destruction of hemoglobin.) This results in a movement of all the protein chains that leads to the other subunits of hemoglobin changing shape to a form with larger oxygen affinity. Thus, when deoxyhemoglobin takes up oxygen, its affinity for more oxygen increases, and vice versa.[141] Myoglobin, on the other hand, contains only one heme group and hence this cooperative effect cannot occur. Thus, while hemoglobin is almost saturated with oxygen in the high partial pressures of oxygen found in the lungs, its affinity for oxygen is much lower than myoglobin in the low partial pressures of oxygen found in muscle tissue, resulting in oxygen transfer.[141] This is further enhanced by the concomitant Bohr effect (named after Christian Bohr, the father of Niels Bohr), in which lowered pH (as occurs when carbon dioxide is released in the muscles) further lowers the oxygen affinity of hemoglobin.[141]

text: Carbon monoxide and phosphorus trifluoride are poisonous to humans because they bind to hemoglobin similarly to oxygen, but with much more strength, so that oxygen can no longer be transported throughout the body. Hemoglobin bound to carbon monoxide is known as carboxyhemoglobin. This effect also plays a minor role in the toxicity of cyanide, but there the major effect is by far its interference with the proper functioning of the electron transport protein cytochrome a.[141] The cytochrome proteins also involve heme groups and are involved in the metabolic oxidation of glucose by oxygen. The sixth coordination site is then occupied by either another imidazole nitrogen or a methionine sulfur, so that these proteins are largely inert to oxygen – with the exception of cytochrome a, which bonds directly to oxygen and thus is very easily poisoned by cyanide.[141] Here, the electron transfer takes place as the iron remains in low spin but changes between the +2 and +3 oxidation states. Since the reduction potential of each step is slightly greater than the previous one, the energy is released step-by-step and can thus be stored in adenosine triphosphate. Cytochrome a is slightly distinct, as it occurs at the mitochondrial membrane, binds directly to oxygen, and transports protons as well as electrons, as follows:[141]

text: Although the heme proteins are the most important class of iron-containing proteins, the iron-sulfur proteins are also very important, being involved in electron transfer, which is possible since iron can exist stably in either the +2 or +3 oxidation states. These have one, two, four, or eight iron atoms that are each approximately tetrahedrally coordinated to four sulfur atoms; because of this tetrahedral coordination, they always have high-spin iron. The simplest of such compounds is rubredoxin, which has only one iron atom coordinated to four sulfur atoms from cysteine residues in the surrounding peptide chains. Another important class of iron-sulfur proteins is the ferredoxins, which have multiple iron atoms. Transferrin does not belong to either of these classes.[141]

text: The ability of sea mussels to maintain their grip on rocks in the ocean is facilitated by their use of organometallic iron-based bonds in their protein-rich cuticles. Based on synthetic replicas, the presence of iron in these structures increased elastic modulus 770 times, tensile strength 58 times, and toughness 92 times. The amount of stress required to permanently damage them increased 76 times.[153]

text: Health and diet

text: Iron is pervasive, but particularly rich sources of dietary iron include red meat, lentils, beans, poultry, fish, leaf vegetables, watercress, tofu, chickpeas, black-eyed peas, and blackstrap molasses. Bread and breakfast cereals are sometimes specifically fortified with iron. Iron in low amounts is found in molasses, teff, and farina.[154][155]

text: Iron provided by dietary supplements is often found as iron(II) fumarate, although iron(II) sulfate is cheaper and is absorbed equally well.[137] Elemental iron, or reduced iron, despite being absorbed at only one-third to two-thirds the efficiency (relative to iron sulfate),[156] 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[157] and is also available for use as a common iron supplement. Glycine, the cheapest and most common amino acid is most often used to produce iron glycinate supplements.[158]

text: Dietary recommendations

text: The U.S. Institute of Medicine (IOM) updated Estimated Average Requirements (EARs) and Recommended Dietary Allowances (RDAs) for iron in 2001. The current EAR for iron for women ages 14–18 is 7.9 mg/day, 8.1 for ages 19–50 and 5.0 thereafter (post menopause). For men the EAR is 6.0 mg/day for ages 19 and up. The RDA is 15.0 mg/day for women ages 15–18, 18.0 for 19-50 and 8.0 thereafter. For men, 8.0 mg/day for ages 19 and up. RDAs are higher than EARs so as to identify amounts that will cover people with higher than average requirements. RDA for pregnancy is 27 mg/day. RDA for lactation is 9 mg/day. For children ages 1–3 years 7 mg/day, 10 for ages 4–8 and 8 for ages 9–13. As for safety, the IOM also sets Tolerable upper intake levels (ULs) for vitamins and minerals when evidence is sufficient. In the case of iron the UL is set at 45 mg/day. Collectively the EARs, RDAs and ULs are referred to as Dietary Reference Intakes.[159]

text: The European Food Safety Authority (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR. AI and UL defined the same as in United States. For women the PRI is 13 mg/day ages 15-17 years, 16 mg/day for women ages 18 and up who are premenopausal and 11 mg/day postmenopausal. For pregnancy and lactation, 16 mg/day. For men the PRI is 11 mg/day ages 15 and older. For children ages 1 to 14 the PRI increases from 7 to 11 mg/day. The PRIs are higher than the U.S. RDAs, with the exception of pregnancy. [160] The EFSA reviewed the same safety question did not establish a UL.[161]

text: Infants may require iron supplements if they are bottle-fed cow's milk.[162] Frequent blood donors are at risk of low iron levels and are often advised to supplement their iron intake.[163]

text: For U.S. food and dietary supplement labeling purposes the amount in a serving is expressed as a percent of Daily Value (%DV). For iron labeling purposes 100% of the Daily Value was 18 mg, and as of May 27, 2016 remained unchanged at 18 mg.[164] A table of all of the old and new adult Daily Values is provided at Reference Daily Intake. The original deadline to be in compliance was July 28, 2018, but on September 29, 2017 the FDA released a proposed rule that extended the deadline to January 1, 2020 for large companies and January 1, 2021 for small companies.[165]

text: Food sources

text: Mild iron deficiency can be prevented or corrected by eating iron-rich foods and by cooking in an iron skillet. Because iron is a requirement for most plants and animals, a wide range of foods provide iron. Good sources of dietary iron have heme-iron, as this is most easily absorbed and is not inhibited by medication or other dietary components. Three examples are red meat, poultry, and insects.[166][167] Non-heme sources do contain iron, though it has reduced bioavailability. Examples are lentils, beans, leafy vegetables, pistachios, tofu, fortified bread, and fortified breakfast cereals.

text: Excess

text:
NFPA 704
fire diamond

1
1
0

Fire diamond hazard sign for metallic iron[168]

text: Iron uptake is tightly regulated by the human body, which has no regulated physiological means of excreting iron. Only small amounts of iron are lost daily due to mucosal and skin epithelial cell sloughing, so control of iron levels is primarily accomplished by regulating uptake.[169] 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 and leads to abnormally low levels of hepcidin, a key regulator of the entry of iron into the circulatory system in mammals.[170] In these people, excessive iron intake can result in iron overload disorders, known medically as hemochromatosis. Many people have an undiagnosed genetic susceptibility to iron overload, and are not aware of a family history of the problem. For this reason, people should not take iron supplements unless they suffer from iron deficiency and have consulted a doctor. Hemochromatosis is estimated to be the cause of 0.3 to 0.8% of all metabolic diseases of Caucasians.[171]

text: Overdoses of ingested iron can cause excessive levels of free iron in the blood. High blood levels of free ferrous iron react with peroxides to produce highly reactive free radicals that can damage DNA, proteins, lipids, and other cellular components. Iron toxicity occurs when the cell contains free iron, which generally occurs when iron levels exceed the availability 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, causing adverse effects that include coma, metabolic acidosis, shock, liver failure, coagulopathy, adult respiratory distress syndrome, long-term organ damage, and even death.[172] Humans experience iron toxicity when the iron exceeds 20 milligrams for every kilogram of body mass; 60 milligrams per kilogram is considered a lethal dose.[173] 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.[173] The Dietary Reference Intake (DRI) sets the Tolerable Upper Intake Level (UL) for adults at 45 mg/day. For children under fourteen years old the UL is 40 mg/day.[174]

text: 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.[172][175][176]

text: Deficiency

text: Iron deficiency is the most common nutritional deficiency in the world.[177][178][179] When loss of iron is not adequately compensated by adequate dietary iron intake, a state of latent iron deficiency occurs, which over time leads to iron-deficiency anemia if left untreated, which is characterised by an insufficient number of red blood cells and an insufficient amount of hemoglobin.[180] Children, pre-menopausal women (women of child-bearing age), and people with poor diet are most susceptible to the disease. Most cases of iron-deficiency anemia are mild, but if not treated can cause problems like fast or irregular heartbeat, complications during pregnancy, and delayed growth in infants and children.[181]


document_title: List of Padma Vibhushan award recipients

question_text: who has been awarded padma vibhushan in 2018

root

text: Recipients[edit]

text:
Award recipients by year[13]

Year

Number of recipients

1954–59

17

1960–69

27

1970–79

53

1980–89

20

1990–99

42

2000–09

86

2010–present

58

Award recipients by field[13]

Field

Number of recipients

Arts

57

Civil Service

53

Literature & Education

40

Medicine

13

Others

5

Public Affairs

68

Science & Engineering

35

Social Work

17

Sports

3

Trade & Industry

12

tag: table

text:
Key

   # Indicates a posthumous honour

text:
List of Padma Vibhushan recipients, showing the year, field, and state/country[13]

Year

Recipient

Field

State

1954

Satyendra Nath Bose

Literature & Education
West Bengal

1954

Nandalal Bose

Arts
West Bengal

1954

Zakir Husain

Public Affairs
Andhra Pradesh

1954

Balasaheb Gangadhar Kher

Public Affairs
Maharashtra

1954

V. K. Krishna Menon

Public Affairs
Kerala

1954

Jigme Dorji Wangchuck

Public Affairs
—[A]

1955

Dhondo Keshav Karve

Literature & Education
Maharashtra

1955

J. R. D. Tata

Trade & Industry
Maharashtra

1956

Fazal Ali

Public Affairs
Bihar

1956

Jankibai Bajaj

Social Work
Madhya Pradesh

1956

Chandulal Madhavlal Trivedi

Public Affairs
Madhya Pradesh

1957

Ghanshyam Das Birla

Trade & Industry
Rajasthan

1957

Sri Prakasa

Public Affairs
Andhra Pradesh

1957

M. C. Setalvad

Public Affairs
Maharashtra

1959

John Mathai

Literature & Education
Kerala

1959

Gaganvihari Lallubhai Mehta

Social Work
Maharashtra

1959

Radhabinod Pal

Public Affairs
West Bengal

1960

Naryana Raghvan Pillai

Public Affairs
Tamil Nadu

1962

H. V. R. Iyengar

Civil Service
Tamil Nadu

1962

Padmaja Naidu

Public Affairs
Andhra Pradesh

1962

Vijaya Lakshmi Pandit

Civil Service
Uttar Pradesh

1963

Suniti Kumar Chatterji

Literature & Education
West Bengal

1963

A. Lakshmanaswami Mudaliar

Medicine
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1963

Hari Vinayak Pataskar

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Maharashtra

1964

Acharya Kakasaheb Kalelkar

Literature & Education
Maharashtra

1964

Gopinath Kaviraj

Literature & Education
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1965

Jayanto Nath Chaudhuri

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West Bengal

1965

Mehdi Nawaz Jung

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Andhra Pradesh

1965

Arjan Singh

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Delhi

1966

Valerian Gracias

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Maharashtra

1967

Chandra Kisan Daphtary

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1967

Hafiz Mohamad Ibrahim

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Andhra Pradesh

1967

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Uttar Pradesh

1967

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Civil Service
Andhra Pradesh

1968

Madhav Shrihari Aney

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Madhya Pradesh

1968

Subrahmanyan Chandrasekhar

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—[B]

1968

Prasanta Chandra Mahalanobis

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Delhi

1968

Kirpal Singh

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Delhi

1968

Kalyan Sundaram

Public Affairs
Delhi

1969

Rajeshwar Dayal

Civil Service
Delhi

1969

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Civil Service
Maharashtra

1969

Har Gobind Khorana

Science & Engineering
—[B]

1969

Mohan Sinha Mehta

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Rajasthan

1969

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Uttar Pradesh

1970

Tara Chand

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1970

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1970

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1970

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1970

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1970

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1971

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1971

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West Bengal

1971

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Maharashtra

1971

Uday Shankar

Arts
Maharashtra

1971

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Goa

1971

B. Sivaraman

Civil Service
Tamil Nadu

1972

P. Balacharya Gajendragadkar

Public Affairs
Maharashtra

1972

Aditya Nath Jha#

Public Affairs
Uttar Pradesh

1972

Pratap Chandra Lal

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Punjab

1972

Sam Manekshaw

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Tamil Nadu

1972

Jivraj Narayan Mehta

Public Affairs
Maharashtra

1972

Sardarilal Mathradas Nanda

Civil Service
Delhi

1972

Ghulam Mohammed Sadiq[i]#

Public Affairs
Jammu & Kashmir

1972

Vikram Sarabhai[ii]#

Science & Engineering
Gujarat

1972

Hormasji Maneckji Seervai

Literature & Education
Maharashtra

1973

Basanti Devi

Civil Service
West Bengal

1973

U. N. Dhebar

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Gujarat

1973

Daulat Singh Kothari

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Delhi

1973

Nellie Sengupta

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West Bengal

1973

Nagendra Singh

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Rajasthan

1973

Thirumalraya Swaminathan

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Tamil Nadu

1974

Niren De

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West Bengal

1974

Benode Behari Mukherjee

Arts
Uttarakhand

1974

V. K. R. V. Rao

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Karnataka

1974

Harish Chandra Sarin

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Delhi

1975

C. D. Deshmukh

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Andhra Pradesh

1975

Durgabai Deshmukh

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Maharashtra

1975

Mary Clubwala Jadhav

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Tamil Nadu

1975

Basanti Dulal Nagchaudhuri

Literature & Education
West Bengal

1975

Raja Ramanna

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Karnataka

1975

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Maharashtra

1975

M. S. Subbulakshmi

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Tamil Nadu

1975

Premlila Vithaldas Thackersey

Literature & Education
Maharashtra

1976

Salim Ali

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Uttar Pradesh

1976

Giani Gurmukh Singh Musafir

Literature & Education
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1976

K. Shankar Pillai

Arts
Delhi

1976

K. R. Ramanathan

Science & Engineering
Kerala

1976

Satyajit Ray

Arts
West Bengal

1976

Kalu Lal Shrimali

Literature & Education
Uttar Pradesh

1976

Bashir Hussain Zaidi

Literature & Education
Delhi

1977

T. Balasaraswati

Arts
Tamil Nadu

1977

Ali Yavar Jung

Public Affairs
Andhra Pradesh

1977

Ajudhia Nath Khosla

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Delhi

1977

Om Prakash Mehra

Civil Service
Punjab

1977

Ajoy Kumar Mukherjee

Public Affairs
West Bengal

1977

Chandeshwar Prasad Narayan Singh

Literature & Education
Delhi

1980

Bismillah Khan

Arts
Uttar Pradesh

1980

Rai Krishnadasa

Civil Service
Uttar Pradesh

1981

Satish Dhawan

Science & Engineering
Karnataka

1981

Ravi Shankar

Arts
Uttar Pradesh

1982

Mira Behn

Social Work
—[C]

1985

C. N. R. Rao

Science & Engineering
Karnataka

1985

M. G. K. Menon

Civil Service
Kerala

1986

Baba Amte[h]

Social Work
Maharashtra

1986

Birju Maharaj

Arts
Delhi

1986

Autar Singh Paintal

Medicine
Delhi

1987

Kamaladevi Chattopadhyay

Social Work
Karnataka

1987

Benjamin Peary Pal

Science & Engineering
Punjab

1987

Manmohan Singh

Civil Service
Delhi

1987

Arun Shridhar Vaidya

Civil Service
Maharashtra

1988

Mirza Hameedullah Beg

Public Affairs
Delhi

1988

Kuvempu

Literature & Education
Karnataka

1988

Mahadevi Varma

Literature & Education
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1989

Uma Shankar Dikshit

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Uttar Pradesh

1989

Ali Akbar Khan

Arts
West Bengal

1989

M. S. Swaminathan

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Delhi

1990

V. S. R. Arunachalam

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Delhi

1990

Triloki Nath Chaturvedi

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Karnataka

1990

Bhabatosh Datta

Literature & Education
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1990

Kumar Gandharva

Arts
Madhya Pradesh

1990

A. P. J. Abdul Kalam

Science & Engineering
Tamil Nadu

1990

Semmangudi Srinivasa Iyer

Arts
Tamil Nadu

1991

M. Balamuralikrishna

Arts
Tamil Nadu

1991

M. F. Husain

Arts
Maharashtra

1991

Hirendranath Mukherjee

Public Affairs
West Bengal

1991

Gulzarilal Nanda

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Gujarat

1991

I. G. Patel

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Gujarat

1991

N. G. Ranga

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Andhra Pradesh

1991

Khusro Faramurz Rustamji

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Maharashtra

1991

Rajaram Shastri

Literature & Education
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1992

Aruna Asaf Ali

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Delhi

1992

Lakshman Shastri Joshi

Literature & Education
Maharashtra

1992

Mallikarjun Mansur

Arts
Karnataka

1992

S. I. Padmavati

Medicine
Delhi

1992

Kaloji Narayana Rao

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Andhra Pradesh

1992

Ravi Narayana Reddy

Public Affairs
Andhra Pradesh

1992

V. Shantaram

Arts
Maharashtra

1992

Govindbhai Shroff

Literature & Education
Maharashtra

1992

Swaran Singh

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Punjab

1992

Atal Bihari Vajpayee

Public Affairs
Delhi

1998

Usha Mehta

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Maharashtra

1998

Nanabhoy Palkhivala

Public Affairs
Maharashtra

1998

Lakshmi Sahgal

Public Affairs
Uttar Pradesh

1998

Walter Sisulu

Public Affairs
—[D]

1999

Pandurang Shastri Athavale

Social Work
Maharashtra

1999

Rajagopala Chidambaram

Science & Engineering
Maharashtra

1999

Nanaji Deshmukh

Social Work
Delhi

1999

Sarvepalli Gopal

Literature & Education
Tamil Nadu

1999

Satish Gujral

Arts
Delhi

1999

V. R. Krishna Iyer

Public Affairs
Kerala

1999

Bhimsen Joshi

Arts
Maharashtra

1999

Hans Raj Khanna

Public Affairs
Delhi

1999

Verghese Kurien

Science & Engineering
Gujarat

1999

Lata Mangeshkar

Arts
Maharashtra

1999

Braj Kumar Nehru

Civil Service
Himachal Pradesh

1999

D. K. Pattammal

Arts
Tamil Nadu

1999

Lallan Prasad Singh[iii]#

Civil Service
Delhi

1999

Dharma Vira

Civil Service
Delhi

2000

Sikander Bakht

Public Affairs
Delhi

2000

Jagdish Bhagwati

Literature & Education
—[B]

2000

Hariprasad Chaurasia

Arts
Maharashtra

2000

M. S. Gill

Civil Service
Delhi

2000

Krishnaswamy Kasturirangan

Science & Engineering
Karnataka

2000

K. B. Lall

Civil Service
Delhi

2000

Kelucharan Mohapatra

Arts
Odisha

2000

Jasraj Motiram

Arts
Maharashtra

2000

M. Narasimham

Trade & Industry
Andhra Pradesh

2000

R. K. Narayan

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Tamil Nadu

2000

Bhairab Dutt Pande

Civil Service
Uttarakhand

2000

K. N. Raj

Literature & Education
Kerala

2000

Tarlok Singh

Civil Service
Delhi

2001

John Kenneth Galbraith

Literature & Education
—[B]

2001

Benjamin A. Gilman

Public Affairs
—[B]

2001

Amjad Ali Khan

Arts
Delhi

2001

Zubin Mehta

Arts
—[B]

2001

Hrishikesh Mukherjee

Arts
Maharashtra

2001

Kotha Satchidanda Murthy

Literature & Education
Andhra Pradesh

2001

Chakravarthi V. Narasimhan

Civil Service
Tamil Nadu

2001

Hosei Norota

Public Affairs
—[E]

2001

C. R. Rao

Science & Engineering
—[B]

2001

Man Mohan Sharma

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Maharashtra

2001

Shivkumar Sharma

Arts
Maharashtra

2002

Kishori Amonkar

Arts
Maharashtra

2002

Gangubai Hangal

Arts
Karnataka

2002

Kishan Maharaj

Arts
Uttar Pradesh

2002

C. Rangarajan

Literature & Education
Andhra Pradesh

2002

Soli Sorabjee

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Delhi

2003

Kazi Lhendup Dorjee

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West Bengal

2003

Sonal Mansingh

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Delhi

2003

Bal Ram Nanda

Literature & Education
Delhi

2003

Brihaspati Dev Triguna

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Delhi

2004

Jayant Narlikar

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Maharashtra

2004

Amrita Pritam

Literature & Education
Delhi

2004

M. N. Venkatachaliah

Public Affairs
Karnataka

2005

Milon K. Banerji

Public Affairs
Delhi

2005

Mohan Dharia

Social Work
Maharashtra

2005

Jyotindra Nath Dixit[iv]#

Civil Service
Delhi

2005

B. K. Goyal

Medicine
Maharashtra

2005

R. K. Laxman

Arts
Maharashtra

2005

Ram Narayan

Arts
Maharashtra

2005

Karan Singh

Public Affairs
Delhi

2005

M. S. Valiathan

Medicine
Karnataka

2006

Norman Borlaug

Science & Engineering
—[B]

2006

Charles Correa

Science & Engineering
Maharashtra

2006

Nirmala Deshpande

Social Work
Delhi

2006

Mahasweta Devi

Literature & Education
West Bengal

2006

Adoor Gopalakrishnan

Arts
Kerala

2006

V. N. Khare

Public Affairs
Uttar Pradesh

2006

C. R. Krishnaswamy Rao

Civil Service
Tamil Nadu

2006

Obaid Siddiqi

Science & Engineering
Karnataka

2006

Prakash Narain Tandon

Medicine
Delhi

2007

P. N. Bhagwati

Public Affairs
Delhi

2007

Naresh Chandra

Civil Service
Delhi

2007

Raja Chelliah

Public Affairs
Tamil Nadu

2007

V. Krishnamurthy

Civil Service
Delhi

2007

Fali Sam Nariman

Public Affairs
Delhi

2007

Raja Rao[v]#

Literature & Education
—[B]

2007

Balu Sankaran

Medicine
Delhi

2007

Khushwant Singh[i]

Literature & Education
Delhi

2007

George Sudarshan

Science & Engineering
—[B]

2007

Narinder Nath Vohra

Civil Service
Haryana

2008

Adarsh Sein Anand

Public Affairs
Uttar Pradesh

2008

Viswanathan Anand

Sports
Tamil Nadu

2008

Asha Bhosle

Arts
Maharashtra

2008

P. N. Dhar

Civil Service
Delhi

2008

Edmund Hillary[vi]#

Sports
—[F]

2008

Lakshmi Mittal

Trade & Industry
—[C]

2008

Pranab Mukherjee

Public Affairs
Delhi

2008

N. R. Narayana Murthy

Trade & Industry
Karnataka

2008

Prithvi Raj Singh Oberoi

Trade & Industry
Delhi

2008

Rajendra K. Pachauri

Science & Engineering
Delhi

2008

E. Sreedharan

Science & Engineering
Delhi

2008

Ratan Tata

Trade & Industry
Maharashtra

2008

Sachin Tendulkar

Sports
Maharashtra

2009

Sunderlal Bahuguna

Others
Uttarakhand

2009

Jasbir Singh Bajaj

Medicine
Punjab

2009

D. P. Chattopadhyaya

Literature & Education
West Bengal

2009

Ashok Sekhar Ganguly

Trade & Industry
Maharashtra

2009

Sister Nirmala

Social Work
West Bengal

2009

Anil Kakodkar

Science & Engineering
Maharashtra

2009

Purshotam Lal

Medicine
Uttar Pradesh

2009

G. Madhavan Nair

Science & Engineering
Karnataka

2009

Govind Narain

Public Affairs
Uttar Pradesh

2009

Chandrika Prasad Srivastava

Civil Service
Maharashtra

2010

Ebrahim Alkazi

Arts
Delhi

2010

Venkatraman Ramakrishnan

Science & Engineering
—[C]

2010

Prathap C. Reddy

Trade & Industry
Tamil Nadu

2010

Y. Venugopal Reddy

Public Affairs
Andhra Pradesh

2010

Zohra Sehgal

Arts
Delhi

2010

Umayalpuram K. Sivaraman

Arts
Tamil Nadu

2011

Montek Singh Ahluwalia

Public Affairs
Delhi

2011

Vijay Kelkar

Public Affairs
Maharashtra

2011

Akhlaqur Rahman Kidwai

Public Affairs
Delhi

2011

O. N. V. Kurup

Literature & Education
Kerala

2011

Sitakant Mahapatra

Literature & Education
Odisha

2011

Brajesh Mishra

Civil Service
Delhi

2011

K. Parasaran

Public Affairs
Delhi

2011

Azim Premji

Trade & Industry
Karnataka

2011

Palle Rama Rao

Science & Engineering
Andhra Pradesh

2011

Akkineni Nageswara Rao

Arts
Andhra Pradesh

2011

Kapila Vatsyayan

Arts
Delhi

2011

Homai Vyarawalla

Arts
Gujarat

2012

Bhupen Hazarika[vii]#

Arts
Assam

2012

Mario Miranda[viii]#

Arts
Goa

2012

T. V. Rajeswar

Civil Service
Delhi

2012

Kantilal Hastimal Sancheti

Medicine
Maharashtra

2012

K. G. Subramanyan

Arts
Gujarat

2013

Raghunath Mohapatra

Arts
Odisha

2013

Roddam Narasimha

Science & Engineering
Karnataka

2013

Yash Pal

Science & Engineering
Uttar Pradesh

2013

S. H. Raza

Arts
Delhi

2014

B. K. S. Iyengar

Others
Maharashtra

2014

Raghunath Anant Mashelkar

Science & Engineering
Maharashtra

2015

L. K. Advani

Public Affairs
Gujarat

2015

Amitabh Bachchan

Arts
Maharashtra

2015

Parkash Singh Badal

Public Affairs
Punjab

2015

Veerendra Heggade

Social Work
Karnataka

2015

Dilip Kumar

Arts
Maharashtra

2015

 Rambhadracharya

Others
Uttar Pradesh

2015

M. R. Srinivasan

Science & Engineering
Tamil Nadu

2015

Kottayan Katankot Venugopal

Public Affairs
Delhi

2015

Karim Al Hussaini Aga Khan

Trade & Industry
—[C][G]

2016

V. K. Aatre

Science & Engineering
Karnataka

2016

Dhirubhai Ambani[ix]#

Trade & Industry
Maharashtra

2016

Girija Devi

Arts
West Bengal

2016

Avinash Dixit

Literature & Education
—[B]

2016

Jagmohan

Public Affairs
Delhi

2016

Yamini Krishnamurthy

Arts
Delhi

2016

Rajinikanth

Arts
Tamil Nadu

2016

Ramoji Rao

Literature & Education
Andhra Pradesh

2016

Sri Sri Ravi Shankar[j]

Others
Karnataka

2016

V. Shanta

Medicine
Tamil Nadu

2017

Murli Manohar Joshi

Public Affairs
Uttar Pradesh

2017

Sunder Lal Patwa[x]#

Public Affairs
Madhya Pradesh

2017

Sharad Pawar

Public Affairs
Maharashtra

2017

Udupi Ramachandra Rao

Science & Engineering
Karnataka

2017

P. A. Sangma[xi]#

Public Affairs
Meghalaya

2017

Jaggi Vasudev

Others
Tamil Nadu

2017

K. J. Yesudas

Arts
Kerala

2018

Ilaiyaraaja

Arts
Tamil Nadu

2018

Ghulam Mustafa Khan

Arts
Maharashtra

2018

P. Parameswaran

Literature & Education
Kerala

text: Explanatory notes[edit]

text:
Jump up ^ Per Article 18 (1) of the Constitution of India: Abolition of titles, "no title, not being a military or academic distinction, shall be conferred by the State".[8][10]

Jump up ^ P. N. Haksar was offered the award in 1973 for, among other services, his crucial diplomatic role in brokering the Indo-Soviet Treaty of Friendship and Cooperation and the Shimla Agreement, but declined as "Accepting an award for work done somehow causes an inexplicable discomfort to me."[14]

Jump up ^ Vilayat Khan refused Padma Shri (1964), Padma Bhushan (1968), and Padma Vibhushan (2000) and stated that "the selection committees were incompetent to judge [his] music".[15][16]

Jump up ^ E. M. S. Namboodiripad, the General Secretary of the Communist Party of India (of the Communist Party of India (Marxist) from 1964) and the first Chief Minister of Kerala (1957-59, 1967-69), declined the award in 1992, as it went against his nature to accept a state honour.[17]

Jump up ^ Swami Ranganathananda declined the award in 2000 as it was conferred to him as an individual and not to the Ramakrishna Mission.[16][18]

Jump up ^ Lakshmi Chand Jain died on 14 November 2010, at the age of 84.[19] His family refused to accept the posthumous honour as Jain was against accepting state honours.[20]

Jump up ^ Sharad Anantrao Joshi's family refused to accept the posthumous honour as Joshi's work for good of farmers is not reflected in the Government policies for them.[21]

^ Jump up to: a b In 1991, Baba Amte returned the award, along with the Padma Shri conferred in 1971, to protest against the treatment given to the tribals during the construction of Sardar Sarovar Dam.[50]

Jump up ^ Khushwant Singh was awarded the Padma Bhushan, the third-highest civilian award,[1] in 1974; he returned the award in 1984 in protest against Operation Blue Star led by the Indian Army.[54]

Jump up ^ Sri Sri Ravi Shankar declined the award in 2015 and requested that "someone else should be given the honour".[59]

tag: ol

text:
Jump up ^ Indicates a citizen of Bhutan

^ Jump up to: a b c d e f g h i j k Indicates a citizen of the United States

^ Jump up to: a b c d Indicates a citizen of the United Kingdom

Jump up ^ Indicates a citizen of South Africa

Jump up ^ Indicates a citizen of Japan

Jump up ^ Indicates a citizen of New Zealand

Jump up ^ Indicates a citizen of France

text:
Jump up ^ Ghulam Mohammed Sadiq died on 12 December 1971, at the age of 59.[48]

Jump up ^ Vikram Sarabhai died on 30 December 1971, at the age of 52.[49]

Jump up ^ Lallan Prasad Singh died on 9 November 1998, at the age of 86.[51]

Jump up ^ Jyotindra Nath Dixit died on 3 January 2005, at the age of 68.[52]

Jump up ^ Raja Rao died on 8 July 2006, at the age of 97.[53]

Jump up ^ Edmund Hillary died on 11 January 2008, at the age of 88.[55]

Jump up ^ Bhupen Hazarika died on 5 November 2011, at the age of 85.[56]

Jump up ^ Mario Miranda died on 11 December 2011, at the age of 85.[57]

Jump up ^ Dhirubhai Ambani died on 6 July 2002, at the age of 69.[58]

Jump up ^ Sunder Lal Patwa died on 28 December 2016, at the age of 92.[60]

Jump up ^ P. A. Sangma died on 4 March 2016, at the age of 68.[61]


document_title: Ceremonial first pitch

question_text: who threw the first pitch at the 2001 world series

text: Ceremonial first pitch

text:
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text: The ceremonial first pitch is a longstanding ritual of baseball in which a guest of honor throws a ball to mark the end of pregame festivities and the start of the game. Originally, the guest threw a ball from his/her place in the grandstand to the pitcher or catcher of the home team, but the ritual changed after Ronald Reagan threw the first pitch on the field at an unscheduled appearance at a Baltimore Orioles game. Now, the guest stands in front of the pitcher's mound and throws towards home plate. He or she may also sometimes stand on the mound (as a pitcher would). The recipient of the pitch is usually a player from the home team.

tag: p

text: The ceremonial thrower may be a notable person (dignitary, celebrity, former player, etc.) who is in attendance, an executive from a company that sponsors the team (especially when that company has sponsored that night's promotional giveaway), or a person who won the first pitch opportunity as a contest prize. Often, especially in the minor leagues, multiple first pitches are made.

text: The practice of having a mayor, governor, or other local celebrity throw out ceremonial "first pitches" dates back to at least 1890.[1] Governor William McKinley of Ohio, for example, "threw the ball into the diamond" before an opening day game between Toledo and Columbus in 1892.[2]

text: On April 23, 2012, the Texas Rangers executed a unique twist on the first pitch tradition. Before the Rangers' home game against the New York Yankees, the team held an official retirement ceremony for longtime catcher Iván Rodríguez. Instead of going to the pitcher's mound, he went behind home plate and threw the first "pitch" to longtime teammate Michael Young, who was standing at second base.[3]

text: Presidential first pitches[edit]

text: Former Japanese Prime Minister Ōkuma Shigenobu threw out the ceremonial first pitch at the first game of an American All-Star team's tour of Japan in 1908, making him possibly the first former chief executive of any country to throw out a first pitch.[1][4]

text: President William Howard Taft started the American tradition of Presidential first pitches in 1910 at Griffith Stadium, Washington, D.C., on the Washington Senators' Opening Day. Every President from Taft to Obama has thrown out at least one ceremonial first ball or pitch, either for Opening Day, the All-Star Game, or the World Series, usually with much fanfare.[5] President Trump has yet to throw a ceremonial first pitch.

text:
Presidential First Pitches

Year
President
Ballpark
Notes

1910 Opening Day
William Howard Taft
National Park
First sitting President to participate in Opening Day ceremonies; preceded Washington Nationals–Philadelphia Athletics game on April 14.[6]

1911 Opening Day
William Howard Taft
National Park
The National Park where the first-ever, Presidential Ceremonial first pitch was thrown burned down in March of 2011, and a new stadium, also called National Park at first, was built in its place. It would be renamed Griffith Stadium in 1923.

1912 Opening Day
Vice President James S. Sherman
National Park
Taft did not attend because of the death of his friend Archibald Butt in the Titanic disaster.

1913 Opening Day
Woodrow Wilson
National Park

1915 Opening Day
Woodrow Wilson
National Park

1915 World Series
Woodrow Wilson
Baker Bowl

1916 Home Opener
Woodrow Wilson
National Park
Washington Nationals defeated New York Yankees on April 20.[7]

1921 Opening Day
Warren Harding
Griffith Stadium
First loss for the Senators with a President throwing out the first ball.

1922 Opening Day
Warren Harding
Griffith Stadium

1923 Opening Day
Warren Harding
Griffith Stadium
Harding also threw out the first ball in the New York Yankees opening game, two days before.

1924 Opening Day
Calvin Coolidge
Griffith Stadium

1924 World Series
Calvin Coolidge
Griffith Stadium

1925 Opening Day
Calvin Coolidge
Griffith Stadium

1925 World Series
Calvin Coolidge
Griffith Stadium

1927 Opening Day
Calvin Coolidge
Griffith Stadium

1928 Opening Day
Calvin Coolidge
Griffith Stadium

1929 Opening Day
Herbert Hoover
Griffith Stadium

1929 World Series
Herbert Hoover
Shibe Park
Game 5, October 14.

1930 Opening Day
Herbert Hoover
Griffith Stadium

1930 World Series
Herbert Hoover
Shibe Park

1931 Opening Day
Herbert Hoover
Griffith Stadium

1932 Opening Day
Herbert Hoover
Griffith Stadium

1933 Opening Day
Franklin D. Roosevelt
Griffith Stadium

1934 Opening Day
Franklin D. Roosevelt
Griffith Stadium

1935 Opening Day
Franklin D. Roosevelt
Griffith Stadium

1936 Opening Day
Franklin D. Roosevelt
Griffith Stadium

1937 Opening Day
Franklin D. Roosevelt
Griffith Stadium

1937 All-Star Game
Franklin D. Roosevelt
Griffith Stadium

1938 Opening Day
Franklin D. Roosevelt
Griffith Stadium

1940 Opening Day
Franklin D. Roosevelt
Griffith Stadium
Roosevelt's pitch hit a Washington Post camera.[8]

1941 Opening Day
Franklin D. Roosevelt
Griffith Stadium

1946 Opening Day
Harry Truman
Griffith Stadium

1947 Opening Day
Harry Truman
Griffith Stadium

1948 Opening Day
Harry Truman
Griffith Stadium

1949 Opening Day
Harry Truman
Griffith Stadium

1950 Opening Day
Harry Truman
Griffith Stadium
Truman threw out two balls, one left-handed and one right-handed.

1951 Opening Day
Harry Truman
Griffith Stadium

1952 Opening Day
Harry Truman
Griffith Stadium

1953 Opening Day
Dwight Eisenhower
Griffith Stadium
Eisenhower skipped Opening Day to play golf at Augusta National, but the game was postponed by rain and he threw out the first ball at the rescheduled game.

1954 Opening Day
Dwight Eisenhower
Griffith Stadium

1955 Opening Day
Dwight Eisenhower
Griffith Stadium

1955 World Series
Dwight Eisenhower
Ebbets Field

1956 Opening Day
Dwight Eisenhower
Griffith Stadium

1958 Opening Day
Dwight Eisenhower
Griffith Stadium

1959 Opening Day
Richard Nixon
Griffith Stadium
Nixon was vice president at the time; he was elected president in 1968.

1960 Opening Day
Dwight Eisenhower
Griffith Stadium

1961 Opening Day
John F. Kennedy
Griffith Stadium

1962 Opening Day
John F. Kennedy
D.C. Stadium
First ball in new stadium

1963 Opening Day
John F. Kennedy
D.C. Stadium

1964 Opening Day
Lyndon Johnson
D.C. Stadium

1965 Opening Day
Lyndon Johnson
D.C. Stadium

1967 Opening Day
Lyndon Johnson
D.C. Stadium

1969 Opening Day
Richard Nixon
RFK Stadium

1970 All-Star Game
Richard Nixon
Riverfront Stadium
All-Star Game in Cincinnati, Ohio

1973 Opening Day
Richard Nixon
Anaheim Stadium
First Opening Day held outside of Washington, D.C.

1976 Opening Day
Gerald Ford
Arlington Stadium

1976 All-Star Game
Gerald Ford
Veterans Stadium
American League vs. National League

1979 World Series
Jimmy Carter
Memorial Stadium

1984 Opening Day
Ronald Reagan
Memorial Stadium
Reagan made an unannounced trip to Baltimore, after it was initially cancelled for security reasons. He watched the game from the third-base dugout.[9]

1986 Opening Day
Ronald Reagan
Memorial Stadium

1988 Regular Season
Ronald Reagan
Wrigley Field
Reagan threw two pitches prior to the Chicago Cubs–Pittsburgh Pirates game on September 30, then joined Harry Caray for 1½ innings on the WGN telecast.[6]

1989 Opening Day
George H. W. Bush
Memorial Stadium
President of Egypt Hosni Mubarak was Bush's special guest, but did not partake in the pregame ceremonies. The Baltimore Orioles defeated the Boston Red Sox 5–4 in 11 innings.[10]

1989 Japan Series
Ronald Reagan
Tokyo Dome
Game 3 between the Kintetsu Buffaloes and Yomiuri Giants

1990 Opening Day
George H. W. Bush
SkyDome
First Opening Day pitch to be thrown in Canada.

1991 Opening Day
George H. W. Bush
Arlington Stadium

1991 Opening Day Ronald Reagan Dodger Stadium

1992 Opening Day
George H. W. Bush
Oriole Park at Camden Yards
First MLB game at Camden Yards.

1993 Opening Day
Bill Clinton
Oriole Park at Camden Yards
Before this, most presidents threw from the stands or at the base of the pitcher's mound; Clinton was the first president to successfully throw from the pitcher's mound to the catcher.

1994 Opening Day
Bill Clinton
Jacobs Field
First MLB Game at Jacobs Field.

1995 World Series
Jimmy Carter
Atlanta–Fulton County Stadium

1996 Opening Day
Bill Clinton
Oriole Park at Camden Yards

1997 Opening Day
Bill Clinton
Shea Stadium

2000 Opening Day
Bill Clinton
Pacific Bell Park

2001 Opening Day
George W. Bush
Miller Park
MLB Commissioner Bud Selig (a former owner of the Brewers) threw out the first pitch to celebrate the opening of the new park; Bush threw the second pitch.

2001 World Series
George W. Bush
Yankee Stadium
This was the first World Series game in New York since the September 11 attacks; Bush wore a bulletproof vest and a Secret Service agent dressed as an umpire so he could be on the field.

2003 Opening Day
George H.W. Bush
Great American Ball Park

2004 Opening Day
Jimmy Carter
Petco Park
First MLB Game at Petco Park

2004 Opening Day
George W. Bush
Busch Stadium

2005 Opening Day
George W. Bush
RFK Stadium
2005 was the Nationals' first season, making Bush the first president to throw out first pitch in Washington since Richard Nixon in 1969.

2006 Opening Day
George W. Bush
Great American Ball Park
First sitting President to participate on Opening Day in Cincinnati; preceded Reds–Cubs game on April 3.[6]

2008 Opening Day
George W. Bush
Nationals Park
This was the first pitch in new stadium. Bush also participated in ESPN's TV broadcast of the game and called the ballpark's first home run, hit by the Braves' Chipper Jones in the 4th inning.

2009 Opening Day
George W. Bush
Rangers Ballpark in Arlington
Bush had owned the Rangers in the early 1990s.

2009 All-Star Game
Barack Obama
Busch Stadium

2009 Japan Series
George W. Bush
Tokyo Dome
Game 3 between the Hokkaido Nippon Ham Fighters and the Yomiuri Giants.

2010 Opening Day
Barack Obama
Nationals Park
100th anniversary of the first Presidential Opening Day ceremonial first pitch

2010 World Series
George W. Bush
Rangers Ballpark in Arlington
Former President Bush was accompanied to the mound by his father, George H.W. Bush, and Texas Rangers team president Nolan Ryan.

2011 College World Series
George W. Bush
TD Ameritrade Park Omaha
This pitch marked the first game at the new home of the College World Series, replacing the nearby Johnny Rosenblatt Stadium. Before Bush threw out the first pitch, his father, who played for Yale in the first CWS in 1947, delivered a video message christening the new stadium. He is the first President to have thrown ceremonial first pitches for amateur and professional (both North America and Japanese) championship matches.[11]

2015 American League Division Series
George H. W. Bush
Minute Maid Park
Former President George H. W. Bush, aged 91, accompanied by his wife Barbara and in a wheelchair with a neck brace, threw the Houston Astros' first pitch at Game 3 of the ALDS against the Kansas City Royals.[12]

2017 World Series
George W. Bush
Minute Maid Park
Former President Bush was accompanied and given the first pitch ball by his father, George H.W. Bush.