Editing Iron (section)

==Chemistry and compounds==
{{Main|Iron compounds}}

{| class="wikitable" style="float:right; clear:right; margin-left:1em; margin-top:0;"
|-
! Oxidation <br />state !! Representative compound
|-
| −2 (d<sup>10</sup>) || [[Disodium tetracarbonylferrate]] (Collman's reagent)
|-
| −1 (d<sup>9</sup>) || {{chem|Fe|2|(CO)|8|2-}}
|-
| 0 (d<sup>8</sup>) || [[Iron pentacarbonyl]]
|-
| 1 (d<sup>7</sup>) || [[Cyclopentadienyliron dicarbonyl dimer]] ("Fp<sub>2</sub>")
|-
| 2 (d<sup>6</sup>) || [[Ferrous sulfate]], [[Ferrocene]]
|-
| 3 (d<sup>5</sup>) || [[Ferric chloride]], [[Ferrocenium tetrafluoroborate]]
|-
| 4 (d<sup>4</sup>) || {{chem|Fe(diars)|2|Cl|2|2+}}, FeO(BF<sub>4</sub>)<sub>2</sub>
|-
| 5 (d<sup>3</sup>) || {{chem|FeO|4|3-}}
|-
| 6 (d<sup>2</sup>) || [[Potassium ferrate]]
|-
|7 (d<sup>1</sup>)
|[FeO<sub>4</sub>]<sup>–</sup> (matrix isolation, 4K)
|}

Iron shows the characteristic chemical properties of the [[transition metal]]s, 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.{{sfn|Greenwood|Earnshaw|1997|p=905}} 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.{{sfn|Greenwood|Earnshaw|1997|p=1070}} Its 26 electrons are arranged in the [[electron configuration|configuration]] [Ar]3d<sup>6</sup>4s<sup>2</sup>, of which the 3d and 4s electrons are relatively close in energy, and thus a number of electrons can be ionized.{{sfn|Greenwood|Earnshaw|1997|pp=1074–75}}

Iron forms compounds mainly in the [[oxidation state]]s +2 ([[iron(II)]], "ferrous") and +3 ([[iron(III)]], "ferric"). Iron also occurs in [[high-valent iron|higher oxidation states]], e.g., the purple [[potassium ferrate]] (K<sub>2</sub>FeO<sub>4</sub>), which contains iron in its +6 oxidation state. The anion [FeO<sub>4</sub>]<sup>–</sup> with iron in its +7 oxidation state, along with an iron(V)-peroxo isomer, has been detected by infrared spectroscopy at 4 K after cocondensation of laser-ablated Fe atoms with a mixture of O<sub>2</sub>/Ar.<ref>{{Cite journal|last1=Lu|first1=Jun-Bo|last2=Jian|first2=Jiwen|last3=Huang|first3=Wei|last4=Lin|first4=Hailu|last5=Li|first5=Jun|last6=Zhou|first6=Mingfei|date=2016-11-16|title=Experimental and theoretical identification of the Fe(VII) oxidation state in FeO<sub>4</sub><sup>−</sup>|journal=Phys. Chem. Chem. Phys.|volume=18|issue=45|pages=31125–31131|doi=10.1039/c6cp06753k|pmid=27812577|bibcode=2016PCCP...1831125L}}</ref> Iron(IV) is a common intermediate in many biochemical oxidation reactions.<ref>{{Cite journal|doi = 10.1021/ar700027f|title = High-Valent Iron(IV)–Oxo Complexes of Heme and Non-Heme Ligands in Oxygenation Reactions|date = 2007|last1 = Nam|first1 = Wonwoo|journal = Accounts of Chemical Research|volume = 40|pages = 522–531|pmid = 17469792|issue = 7|url = https://cbs.ewha.ac.kr/pub/data/2007_07.pdf|access-date = 22 February 2022|archive-date = 15 June 2021|archive-url = https://web.archive.org/web/20210615123946/http://cbs.ewha.ac.kr/pub/data/2007_07.pdf|url-status = dead}}</ref><ref name="HollemanAF">{{Cite book|publisher = Walter de Gruyter|date = 1985|edition = 91–100|pages = 1125–46|isbn = 3-11-007511-3|title = Lehrbuch der Anorganischen Chemie|first1 = Arnold F.|last1 = Holleman|last2 = Wiberg|first2 = Egon|last3 = Wiberg|first3 = Nils|chapter = Iron| language = de}}</ref> Numerous [[organoiron chemistry|organoiron]] 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]].<ref>{{Cite book| chapter = Mössbauer Spectroscopy and the Coordination Chemistry of Iron|first1 = William Michael|last1 = Reiff|first2 = Gary J.|last2 = Long |title = Mössbauer spectroscopy applied to inorganic chemistry|publisher = Springer|date = 1984|isbn = 978-0-306-41647-7|pages = 245–83}}</ref> Many [[mixed valence compound]]s contain both iron(II) and iron(III) centers, such as [[magnetite]] and [[Prussian blue]] ({{chem2|Fe4(Fe[CN]6)3}}).<ref name="HollemanAF" /> The latter is used as the traditional "blue" in [[blueprint]]s.<ref>{{Cite book| chapter = An introduction in monochrome|pages = 11–19|first = Mike|last = Ware|publisher = NMSI Trading Ltd|title = Cyanotype: the history, science and art of photographic printing in Prussian blue|isbn = 978-1-900747-07-3|date = 1999| chapter-url={{Google books|C-7I69gFIbMC|page=PA11|keywords=|text=|plainurl=yes}}}}</ref>

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.{{sfn|Greenwood|Earnshaw|1997|pp=1075–79}} 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.{{sfn|Greenwood|Earnshaw|1997|pp=1075–79}} In the 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]].{{sfn|Greenwood|Earnshaw|1997|p=1070}}

Unlike many other metals, iron does not form amalgams with [[mercury (element)|mercury]]. As a result, mercury is traded in standardized 76 pound flasks (34 kg) made of iron.<ref>{{Cite book|title = Hand-book of chemistry|volume = 6| first1 = Leopold|last1 = Gmelin|author-link = Leopold Gmelin|pages = 128–29| chapter = Mercury and Iron|chapter-url={{Google books|nosMAAAAYAAJ|page=PA128|keywords=|text=|plainurl=yes}}|publisher = Cavendish Society|date = 1852}}</ref>

Iron is by far the most reactive element in its group; it is [[pyrophoricity|pyrophoric]] when finely divided and dissolves easily in dilute acids, giving Fe<sup>2+</sup>. 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]].{{sfn|Greenwood|Earnshaw|1997|pp=1075–79}} High-purity iron, called [[electrolytic iron]], is considered to be resistant to rust, due to its oxide layer.

===Binary compounds===
====Oxides and sulfides====
{{multiple image
| align = right
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| width = 160
| image1 = Iron(II) oxide.jpg
| caption1 = Ferrous or iron(II) oxide, {{chem2|FeO}}
| image2 = Iron(III)-oxide-sample.jpg
| caption2 = Ferric or iron(III) oxide {{chem2|Fe2O3}}
| image3 = Fe3O4.JPG
| caption3 = Ferrosoferric or iron(II,III) oxide {{chem2|Fe3O4}}
| total_width = 
| alt1 = 
}}
Iron forms various [[iron oxide|oxide and hydroxide compounds]]; the most common are [[iron(II,III) oxide]] (Fe<sub>3</sub>O<sub>4</sub>), and [[iron(III) oxide]] (Fe<sub>2</sub>O<sub>3</sub>). [[Iron(II) oxide]] also exists, though it is unstable at room temperature. Despite their names, they are actually all [[non-stoichiometric compound]]s whose compositions may vary.{{sfn|Greenwood|Earnshaw|1997|p=1079}} These oxides are the principal ores for the production of iron (see [[bloomery]] and blast furnace). They are also used in the production of [[Ferrite (magnet)|ferrites]], useful [[magnetic storage]] media in computers, and pigments. The best known sulfide is [[iron pyrite]] (FeS<sub>2</sub>), also known as fool's gold owing to its golden luster.<ref name="HollemanAF" /> It is not an iron(IV) compound, but is actually an iron(II) [[polysulfide]] containing Fe<sup>2+</sup> and {{chem|S|2|2-}} ions in a distorted [[sodium chloride]] structure.{{sfn|Greenwood|Earnshaw|1997|p=1079}}

[[File:Pourbaix Diagram of Iron.svg|thumb|right|[[Pourbaix diagram]] of iron]]

====Halides====
[[File:Iron(III) chloride hexahydrate.jpg|thumb|upright=0.7|alt=Some canary-yellow powder sits, mostly in lumps, on a laboratory watch glass.|Hydrated iron(III) chloride (ferric chloride)]]
The binary ferrous and ferric [[halide]]s are well-known. The ferrous halides typically arise from treating iron metal with the corresponding [[hydrohalic acid]] to give the corresponding hydrated salts.<ref name="HollemanAF" />
:Fe + 2 HX → FeX<sub>2</sub> + H<sub>2</sub> (X = F, Cl, Br, I)
Iron reacts with fluorine, chlorine, and bromine to give the corresponding ferric halides, [[ferric chloride]] being the most common.{{sfn|Greenwood|Earnshaw|1997|pp=1082–84}}
:2 Fe + 3 X<sub>2</sub> → 2 FeX<sub>3</sub> (X = F, Cl, Br)
[[Ferric iodide]] is an exception, being thermodynamically unstable due to the oxidizing power of Fe<sup>3+</sup> and the high reducing power of I<sup>−</sup>:{{sfn|Greenwood|Earnshaw|1997|pp=1082–84}}
:2 I<sup>−</sup> + 2 Fe<sup>3+</sup> → I<sub>2</sub> + 2 Fe<sup>2+</sup> (E<sup>0</sup> = +0.23 V)

Ferric iodide, a black solid, is not stable in ordinary conditions, but can 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, with oxygen and water excluded.{{sfn|Greenwood|Earnshaw|1997|pp=1082–84}} Complexes of ferric iodide with some soft bases are known to be stable compounds.<ref>Siegfried Pohl, Ulrich Bierbach, Wolfgang Saak; "FeI3SC(NMe2)2, a Neutral Thiourea Complex of Iron(III) Iodide", Angewandte Chemie International Edition in English (1989) 28 (6), 776–777. https://doi.org/10.1002/anie.198907761</ref><ref>Nicholas A. Barnes, Stephen M.Godfrey, Nicholas Ho, Charles A.McAuliffe, Robin G.Pritchard; "Facile synthesis of a rare example of an iron(III) iodide complex, [FeI3(AsMe3)2], from the reaction of Me3AsI2 with unactivated iron powder", Polyhedron (2013) 55, 67–72. https://doi.org/10.1016/j.poly.2013.02.066</ref>

===Solution chemistry===
[[File:Ferrate_and_permanganate_solution.jpg|thumb|upright=0.7|right|Comparison of colors of solutions of ferrate (left) and [[permanganate]] (right)]]
The [[standard reduction potential]]s in acidic aqueous solution for some common iron ions are given below:{{sfn|Greenwood|Earnshaw|1997|pp=1075–79}}
:{|
|-
| [Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup> + 2 e<sup>−</sup>|| {{eqm}} Fe || E<sup>0</sup> = −0.447 V
|-
| [Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>3+</sup> + e<sup>−</sup>|| {{eqm}} [Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup> || E<sup>0</sup> = +0.77 V
|-
| {{chem|FeO|4|2-}} + 8 H<sub>3</sub>O<sup>+</sup> + 3 e<sup>−</sup>|| {{eqm}} [Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>3+</sup> + 6 H<sub>2</sub>O || E<sup>0</sup> = +2.20 V
|}

The red-purple tetrahedral [[ferrate]](VI) anion is such a strong oxidizing agent that it oxidizes ammonia to nitrogen (N<sub>2</sub>) and water to oxygen:{{sfn|Greenwood|Earnshaw|1997|pp=1082–84}}
:4 {{chem|FeO|4|2-}} + 34 {{chem|H|2|O}} → 4 {{chem2|[Fe(H2O)6](3+)}} + 20 {{chem|OH|-}} + 3 O<sub>2</sub>

The pale-violet hex[[aquo complex]] {{chem2|[Fe(H2O)6](3+)}} is an acid such that above pH 0 it is fully hydrolyzed:{{sfn|Greenwood|Earnshaw|1997|pp=1088–91}}

:{|
|-
| {{chem2|[Fe(H2O)6](3+)}} || {{eqm}} {{chem2|[Fe(H2O)5(OH)](2+) + H(+)}} || ''[[equilibrium constant|K]]'' = 10<sup>−3.05</sup> mol dm<sup>−3</sup>
|-
| {{chem2|[Fe(H2O)5(OH)](2+)}} || {{eqm}} {{chem2|[Fe(H2O)4(OH)2](+) + H(+)}} || ''K'' = 10<sup>−3.26</sup> mol dm<sup>−3</sup>
|-
| {{chem2|2[Fe(H2O)6](3+)}} || {{eqm}} {{chem2|[Fe(H2O)4(OH)]2(4+) + 2H(+) + 2H2O}} || ''K'' = 10<sup>−2.91</sup> mol dm<sup>−3</sup>
|}

[[File:Iron(II)-sulfate-heptahydrate-sample.jpg|thumb|right|upright=0.7|Blue-green [[iron(II) sulfate]] heptahydrate]]

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 Fe<sup>3+</sup> has a d<sup>5</sup> configuration, its absorption spectrum is not like that of Mn<sup>2+</sup> with its weak, spin-forbidden d–d bands, because Fe<sup>3+</sup> has higher positive charge and is more polarizing, lowering the energy of its ligand-to-metal [[charge-transfer complex|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.{{sfn|Greenwood|Earnshaw|1997|pp=1088–91}} On the other hand, the pale green iron(II) hexaquo ion {{chem2|[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.{{sfn|Greenwood|Earnshaw|1997|pp=1091–97}}

===Coordination compounds===
Due to its electronic structure, iron has a very large coordination and organometallic chemistry.
[[File:2-isomers-of-ferrioxalate.svg|thumb|right|The two [[enantiomer|enantiomorphs]] of the ferrioxalate ion]]
Many coordination compounds of iron are known. A typical six-coordinate anion is hexachloroferrate(III), [FeCl<sub>6</sub>]<sup>3−</sup>, found in the mixed [[salt (chemistry)|salt]] [[tetrakis(methylammonium) hexachloroferrate(III) chloride]].<ref>{{cite journal|last1 = Clausen|first1 = C.A.|last2 = Good|first2 = M.L.|year = 1968|title = Stabilization of the hexachloroferrate(III) anion by the methylammonium cation|journal = [[Inorganic Chemistry (journal)|Inorganic Chemistry]] |volume = 7|issue = 12|pages = 2662–63|doi = 10.1021/ic50070a047}}</ref><ref>{{cite journal|last1 = James|first1 = B.D.|first2 = M.|last2 = Bakalova|first3 = J.|last3 = Lieseganga|first4 = W.M.|last4 = Reiff|first5 = D.C.R.|last5 = Hockless|first6 = B.W.|last6 = Skelton|first7 = A.H.|last7 = White|year = 1996|title = The hexachloroferrate(III) anion stabilized in hydrogen bonded packing arrangements. A comparison of the X-ray crystal structures and low temperature magnetism of tetrakis(methylammonium) hexachloroferrate(III) chloride '''(I)''' and tetrakis(hexamethylenediammonium) hexachloroferrate(III) tetrachloroferrate(III) tetrachloride '''(II)'''|journal = [[Inorganica Chimica Acta]]|volume = 247|issue = 2|pages = 169–74|doi = 10.1016/0020-1693(95)04955-X}}</ref> Complexes with multiple bidentate ligands have [[geometric isomer]]s. For example, the ''trans''-[[chlorohydridobis(bis-1,2-(diphenylphosphino)ethane)iron(II)]] complex is used as a starting material for compounds with the {{chem2|Fe([[dppe]])2}} [[moiety (chemistry)|moiety]].<ref>{{cite book|last1 = Giannoccaro|first1 = P.|last2 = Sacco|first2 = A.| title=Inorganic Syntheses | chapter=Bis[Ethylenebis(Diphenylphosphine)]-Hydridoiron Complexes|year = 1977|volume = 17|pages = 69–72|doi = 10.1002/9780470132487.ch19|isbn = 978-0-470-13248-7}}</ref><ref>{{cite journal|last1 = Lee|first1 = J.|last2 = Jung|first2 = G.|last3 = Lee|first3 = S.W.|title = Structure of trans-chlorohydridobis(diphenylphosphinoethane)iron(II)|journal = Bull. Korean Chem. Soc.|year = 1998|volume = 19|issue = 2|pages = 267–69|url = https://www.koreascience.or.kr/article/ArticleFullRecord.jsp?cn=JCGMCS_1998_v19n2_267|doi = 10.1007/BF02698412|s2cid = 35665289}}</ref> The ferrioxalate ion with three [[oxalate]] ligands 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.{{sfn|Greenwood|Earnshaw|1997|pp=1088–91}} [[Potassium ferrioxalate]] is used in chemical [[actinometry]] and along with its [[sodium ferrioxalate|sodium salt]] undergoes [[photoreduction]] applied in old-style photographic processes. The [[hydrate|dihydrate]] of [[iron(II) oxalate]] has a [[polymer]]ic 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.<ref>{{cite journal|first1 = Takuya|last1 = Echigo|first2 = Mitsuyoshi|last2 = Kimata|title = Single-crystal X-ray diffraction and spectroscopic studies on humboldtine and lindbergite: weak Jahn–Teller effect of Fe<sup>2+</sup> ion|journal = [[Physics and Chemistry of Minerals|Phys. Chem. Miner.]]|year = 2008|volume = 35|issue = 8|pages = 467–75|doi= 10.1007/s00269-008-0241-7|bibcode = 2008PCM....35..467E|s2cid = 98739882}}</ref>
[[File:Fe(C2O4)(H2O)2-chain-from-xtal-2008-CM-3D-balls.png|upright=2.7|thumb|center|Crystal structure of iron(II) oxalate dihydrate, showing iron (gray), oxygen (red), carbon (black), and hydrogen (white) atoms.]]

[[File:Pentaaqua(thiocyanato)iron(III) chloride.jpg|thumb|upright=0.7|right|Blood-red positive thiocyanate test for iron(III)]]
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 [[phenol]]s or [[enol]]s. 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:{{sfn|Greenwood|Earnshaw|1997|pp=1088–91}}
:3 ArOH + FeCl<sub>3</sub> → Fe(OAr)<sub>3</sub> + 3 HCl (Ar = [[aryl]])

Among the halide and pseudohalide complexes, fluoro complexes of iron(III) are the most stable, with the colorless [FeF<sub>5</sub>(H<sub>2</sub>O)]<sup>2−</sup> being the most stable in aqueous solution. Chloro complexes are less stable and favor tetrahedral coordination as in [FeCl<sub>4</sub>]<sup>−</sup>; [FeBr<sub>4</sub>]<sup>−</sup> and [FeI<sub>4</sub>]<sup>−</sup> are reduced easily to iron(II). [[Thiocyanate]] is a common test for the presence of iron(III) as it forms the blood-red [Fe(SCN)(H<sub>2</sub>O)<sub>5</sub>]<sup>2+</sup>. 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)<sub>6</sub>]<sup>3−</sup>. Iron shows a great variety of electronic [[spin states (d electrons)|spin states]], including every possible spin quantum number value for a d-block element from 0 (diamagnetic) to {{frac|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.{{sfn|Greenwood|Earnshaw|1997|p=1079}}

Iron(II) complexes are less stable than iron(III) complexes but the preference for ''O''-donor ligands is less marked, so that for example {{chem2|[Fe(NH3)6](2+)}} is known while {{chem2|[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.{{sfn|Greenwood|Earnshaw|1997|pp=1091–97}}

===Organometallic compounds===
[[File:Sample of iron pentacarbonyl.jpg|thumb|upright=0.36|left|Iron penta-<br/>carbonyl]]
[[Organoiron chemistry]] is the study of [[organometallic compound]]s of iron, where carbon atoms are covalently bound to the metal atom. They are many and varied, including [[cyanometallate|cyanide complexes]], [[carbonyl complex]]es, [[sandwich compound|sandwich]] and [[half-sandwich compound]]s.

[[File:Prussian blue.jpg|thumb|right|Prussian blue]]
[[Prussian blue]] or "ferric ferrocyanide", Fe<sub>4</sub>[Fe(CN)<sub>6</sub>]<sub>3</sub>, is an old and well-known iron-cyanide complex, extensively used as pigment and in several other applications. Its formation can be used as a simple wet chemistry test to distinguish between aqueous solutions of Fe<sup>2+</sup> and Fe<sup>3+</sup> as they react (respectively) with [[potassium ferricyanide]] and [[potassium ferrocyanide]] to form Prussian blue.<ref name="HollemanAF" />

Another old example of an organoiron compound is [[iron pentacarbonyl]], Fe(CO)<sub>5</sub>, in which a neutral iron atom is bound to the carbon atoms of five [[carbon monoxide]] molecules. The compound can be used to make [[carbonyl iron]] powder, a highly reactive form of metallic iron. [[Thermal decomposition|Thermolysis]] of iron pentacarbonyl gives [[triiron dodecacarbonyl]], {{chem2|Fe3(CO)12}}, a complex with a cluster of three iron atoms at its core. 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.<ref>{{Greenwood&Earnshaw1st|pages=1282–86}}.</ref>

{{multiple image
| align = top | direction = h
| total_width = 250
| image1 = Ferrocene.svg
| image2 = Photo of Ferrocene (powdered).JPG
| footer = Structural formula of ferrocene and a powdered sample
}}
A landmark in this field was the discovery in 1951 of the remarkably stable [[sandwich compound]] [[ferrocene]] {{chem2|Fe(C5H5)2}}, by Pauson and Kealy<ref>{{Cite journal |last1= Kealy |first1=T.J. |last2= Pauson |first2=P.L. |year= 1951 |title= A New Type of Organo-Iron Compound |journal= [[Nature (journal)|Nature]] |volume= 168 |issue= 4285 |pages= 1039–40 |doi= 10.1038/1681039b0|bibcode = 1951Natur.168.1039K |s2cid=4181383 }}</ref> and independently by Miller and colleagues,<ref name = Miller>{{cite journal| last1=Miller|first1= S. A.|last2=Tebboth|first2= J. A.|last3= Tremaine|first3= J. F.|journal= [[Journal of the Chemical Society|J. Chem. Soc.]]|year=1952| pages= 632–635| title=114. Dicyclopentadienyliron |doi=10.1039/JR9520000632}}</ref> whose surprising molecular structure was determined only a year later by [[Robert Burns Woodward|Woodward]] and [[Geoffrey Wilkinson|Wilkinson]]<ref>{{cite journal |first1= G. |last1=Wilkinson|author1-link=Geoffrey Wilkinson |first2=M.|last2= Rosenblum |first3=M. C.|last3= Whiting|first4= R. B. |last4=Woodward|author4-link=Robert Burns Woodward |title = The Structure of Iron Bis-Cyclopentadienyl |journal = [[Journal of the American Chemical Society|J. Am. Chem. Soc.]] |year = 1952|volume = 74 |pages = 2125–2126 |doi = 10.1021/ja01128a527 |issue = 8|bibcode=1952JAChS..74.2125W }}</ref> and [[Ernst Otto Fischer|Fischer]].<ref>{{Cite journal|last=Okuda|first=Jun|date=2016-12-28|title=Ferrocene – 65 Years After|journal=European Journal of Inorganic Chemistry|volume=2017|issue=2|pages=217–219|doi=10.1002/ejic.201601323|issn=1434-1948|doi-access=free}}</ref>
Ferrocene is still one of the most important tools and models in this class.{{sfn|Greenwood|Earnshaw|1997|p=1104}}

Iron-centered organometallic species are used as [[catalyst]]s. The [[Knölker complex]], for example, is a [[transfer hydrogenation]] catalyst for [[ketone]]s.<ref>{{cite journal|last=Bullock|first=R.M.|date=11 September 2007|title=An Iron Catalyst for Ketone Hydrogenations under Mild Conditions|journal=[[Angew. Chem. Int. Ed.]]|volume=46|issue=39|pages=7360–63|doi=10.1002/anie.200703053|pmid=17847139|doi-access=free}}</ref>

===Industrial uses===
The iron compounds produced on the largest scale in industry are [[iron(II) sulfate]] (FeSO<sub>4</sub>·7[[Water of crystallization|H<sub>2</sub>O]]) and [[iron(III) chloride]] (FeCl<sub>3</sub>). The former is one of the most readily available sources of iron(II), but is less stable to aerial oxidation than [[Mohr's salt]] ({{chem2|(NH4)2Fe(SO4)2*6H2O}}). Iron(II) compounds tend to be oxidized to iron(III) compounds in the air.<ref name="HollemanAF" />