View source for Chemistry:Iron compounds

{{Short description|Chemical compounds containing iron}}
{{category see also|Iron compounds}}
{| class="wikitable" style="float:right; clear:right; margin-left:1em; margin-top:0;"
|-
! Oxidation <br />state !! Representative compound
|-
| −4 (d<sup>10</sup>s<sup>2</sup>)|| [FeIn<sub>6−''x''</sub>Sn<sub>''x''</sub>]<ref name="MetalAnions">Fe(−4), Ru(−4), and Os(−4) have been observed in metal-rich compounds containing octahedral complexes [MIn<sub>6−''x''</sub>Sn<sub>''x''</sub>]; Pt(−3) (as a dimeric anion [Pt–Pt]<sup>6−</sup>), Cu(−2), Zn(−2), Ag(−2), Cd(−2), Au(−2), and Hg(−2) have been observed (as dimeric and monomeric anions; dimeric ions were initially reported to be [T–T]<sup>2−</sup> for Zn, Cd, Hg, but later shown to be [T–T]<sup>4−</sup> for all these elements) in La<sub>2</sub>Pt<sub>2</sub>In, La<sub>2</sub>Cu<sub>2</sub>In, Ca<sub>5</sub>Au<sub>3</sub>, Ca<sub>5</sub>Ag<sub>3</sub>, Ca<sub>5</sub>Hg<sub>3</sub>, Sr<sub>5</sub>Cd<sub>3</sub>, Ca<sub>5</sub>Zn<sub>3</sub>(structure (AE<sup>2+</sup>)<sub>5</sub>(T–T)<sup>4−</sup>T<sup>2−</sup>⋅4e<sup>−</sup>), Yb<sub>3</sub>Ag<sub>2</sub>, Ca<sub>5</sub>Au<sub>4</sub>, and Ca<sub>3</sub>Hg<sub>2</sub>; Au(–3) has been observed in ScAuSn and in other 18-electron half-Heusler compounds. See {{cite journal|title=Late transition metal anions acting as p-metal elements|year=2008|author1=Changhoon Lee|author2=Myung-Hwan Whangbo|volume=10|issue=4|pages=444–449|journal=Solid State Sciences|doi=10.1016/j.solidstatesciences.2007.12.001|bibcode=2008SSSci..10..444K}} and {{cite journal|doi=10.1002/zaac.200900421|title=Analysis of Electronic Structures and Chemical Bonding of Metal-rich Compounds. 2. Presence of Dimer (T–T)<sup>4–</sup> and Isolated T<sup>2–</sup> Anions in the Polar Intermetallic Cr<sub>5</sub>B<sub>3</sub>-Type Compounds AE<sub>5</sub>T<sub>3</sub> (AE = Ca, Sr; T = Au, Ag, Hg, Cd, Zn)|year=2010|author1=Changhoon Lee|author2=Myung-Hwan Whangbo|author3=Jürgen Köhler|volume=636|issue=1|pages=36–40|journal=Zeitschrift für Anorganische und Allgemeine Chemie}}</ref>
|-
| −2 (d<sup>10</sup>) || [[Chemistry:Disodium tetracarbonylferrate|Disodium tetracarbonylferrate]] (Collman's reagent)
|-
| −1 (d<sup>9</sup>) || {{chem|Fe|2|(CO)|8|2-}}
|-
| 0 (d<sup>8</sup>) || [[Chemistry:Iron pentacarbonyl|Iron pentacarbonyl]]
|-
| 1 (d<sup>7</sup>) || [[Chemistry:Cyclopentadienyliron dicarbonyl dimer|Cyclopentadienyliron dicarbonyl dimer]] ("Fp<sub>2</sub>")
|-
| 2 (d<sup>6</sup>) || [[Chemistry:Ferrous sulfate|Ferrous sulfate]], [[Chemistry:Ferrocene|ferrocene]]
|-
| 3 (d<sup>5</sup>) || Ferric chloride, [[Chemistry:Ferrocenium tetrafluoroborate|ferrocenium tetrafluoroborate]]
|-
| 4 (d<sup>4</sup>) || {{chem|Fe(diars)|2|Cl|2|2+}}, [[Chemistry:Iron(IV) oxytetrafluoroborate|Ferryl tetrafluoroborate]]
|-
| 5 (d<sup>3</sup>) || {{chem|FeO|4|3-}}
|-
| 6 (d<sup>2</sup>) || [[Chemistry:Potassium ferrate|Potassium ferrate]]
|-
|7 (d<sup>1</sup>)
|[FeO<sub>4</sub>]<sup>–</sup> (matrix isolation, 4K)
|}

Iron shows the characteristic chemical properties of the [[Chemistry:Transition metal|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, [[Chemistry:Ferrocene|ferrocene]], that revolutionalized the latter field in the 1950s.<ref name="Greenwood905">Greenwood and Earnshaw, p. 905</ref> 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.<ref name="Greenwood1070" /> Its 26 electrons are arranged in the [[Physics: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 it can lose a variable number of electrons and there is no clear point where further ionization becomes unprofitable.<ref>{{Greenwood&Earnshaw2nd}} (p. 1074).</ref>

Iron forms compounds mainly in the [[Chemistry:Oxidation state|oxidation state]]s +2 ([[Chemistry:Iron(II)|iron(II)]], "ferrous") and +3 ([[Chemistry:Iron(III)|iron(III)]], "ferric"). Iron also occurs in [[Chemistry:High-valent iron|higher oxidation states]], e.g. the purple [[Chemistry:Potassium ferrate|potassium ferrate]] (K<sub>2</sub>FeO<sub>4</sub>), which contains iron in its +6 oxidation state. Although iron(VIII) oxide (FeO<sub>4</sub>) has been claimed, the report could not be reproduced and such a species from the removal of all electrons of the element beyond the preceding inert gas configuration (at least with iron in its +8 oxidation state) has been found to be improbable computationally.<ref>{{Cite journal|last1=Huang|first1=Wei|last2=Xu|first2=Wen-Hua |last3=Schwarz|first3=W.H.E.|last4=Li|first4=Jun|date=2016-05-02|title=On the Highest Oxidation States of Metal Elements in MO<sub>4</sub> Molecules (M = Fe, Ru, Os, Hs, Sm, and Pu)|journal=Inorganic Chemistry|volume=55|issue=9|pages=4616–25|doi=10.1021/acs.inorgchem.6b00442|pmid=27074099}}</ref> However, one form of anionic [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 [[Chemistry: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 compounds contain both iron(II) and iron(III) centers, such as [[Chemistry:Magnetite|magnetite]] and [[Chemistry:Prussian blue|Prussian blue]] ({{chem2|Fe4(Fe[CN]6)3}}).<ref name="HollemanAF" /> The latter is used as the traditional "blue" in [[Blueprint|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.<ref name="Greenwood1075" /> 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.<ref name="Greenwood1075" /> In the second half of the 3d transition series, vertical similarities down the groups compete with the horizontal similarities of iron with its neighbors [[Chemistry:Cobalt|cobalt]] and [[Chemistry:Nickel|nickel]] in the periodic table, which are also ferromagnetic at [[Physics:Room temperature|room temperature]] and share similar chemistry. As such, iron, cobalt, and nickel are sometimes grouped together as the iron triad.<ref name="Greenwood1070">Greenwood and Earnshaw, p. 1070</ref>

Unlike many other metals, iron does not form amalgams with [[Chemistry:Mercury (element)|mercury]]. As a result, mercury is traded in standardized 76 pound flasks (34&nbsp;kg) made of iron.<ref>{{Cite book|title = Hand-book of chemistry|volume = 6| first1 = Leopold|last1 = 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 [[Chemistry:Pyrophoricity|pyrophoric]] when finely divided and dissolves easily in dilute acids, giving Fe<sup>2+</sup>. However, it does not react with concentrated [[Chemistry:Nitric acid|nitric acid]] and other oxidizing acids due to the formation of an impervious oxide layer, which can nevertheless react with [[Chemistry:Hydrochloric acid|hydrochloric acid]].<ref name="Greenwood1075" /> High purity iron, called [[Physics:Electrolytic iron|electrolytic iron]], is considered to be resistant to rust, due to its oxide layer.

==Binary compounds==
===Oxides and hydroxides===
{{multiple image
| align = right
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| 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 [[Chemistry:Iron oxide|oxide and hydroxide compounds]]; the most common are [[Chemistry:Iron(II,III) oxide|iron(II,III) oxide]] (Fe<sub>3</sub>O<sub>4</sub>), and [[Chemistry:Iron(III) oxide|iron(III) oxide]] (Fe<sub>2</sub>O<sub>3</sub>). [[Chemistry:Iron(II) oxide|Iron(II) oxide]] also exists, though it is unstable at room temperature. Despite their names, they are actually all [[Chemistry:Non-stoichiometric compound|non-stoichiometric compound]]s whose compositions may vary.<ref name="Greenwood1079" /> These oxides are the principal ores for the production of iron (see [[Chemistry:Bloomery|bloomery]] and blast furnace). They are also used in the production of [[Physics:Ferrite (magnet)|ferrites]], useful [[Magnetic storage|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) [[Chemistry:Polysulfide|polysulfide]] containing Fe<sup>2+</sup> and {{chem|S|2|2-}} ions in a distorted [[Chemistry:Sodium chloride|sodium chloride]] structure.<ref name="Greenwood1079">Greenwood and Earnshaw, p. 1079</ref>

[[File:Pourbaix Diagram of Iron.svg|thumb|right|[[Chemistry:Pourbaix diagram|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 [[Chemistry:Halide|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.<ref name="Greenwood1082" />
: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>:<ref name="Greenwood1082" />
: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 [[Chemistry:Iron pentacarbonyl|iron pentacarbonyl]] with [[Chemistry:Iodine|iodine]] and [[Chemistry:Carbon monoxide|carbon monoxide]] in the presence of [[Chemistry:Hexane|hexane]] and light at the temperature of −20&nbsp;°C, with oxygen and water excluded.<ref name="Greenwood1082" /> 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 [[Chemistry:Permanganate|permanganate]] (right)]]
The standard reduction potentials in acidic aqueous solution for some common iron ions are given below:<ref name="Greenwood1075">Greenwood and Earnshaw, pp. 1075–79</ref>
{|
|-
| Fe<sup>2+</sup> + 2 e<sup>−</sup>|| {{eqm}} Fe || E<sup>0</sup> = −0.447 V
|-
| Fe<sup>3+</sup> + 3 e<sup>−</sup>|| {{eqm}} Fe || E<sup>0</sup> = −0.037 V
|-
| {{chem|FeO|4|2-}} + 8 H<sup>+</sup> + 3 e<sup>−</sup>|| {{eqm}} Fe<sup>3+</sup> + 4 H<sub>2</sub>O || E<sup>0</sup> = +2.20 V
|}

The red-purple tetrahedral [[Chemistry:Ferrate|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:<ref name="Greenwood1082">Greenwood and Earnshaw, pp. 1082–84</ref>
:4 {{chem|FeO|4|2-}} + 10 {{chem|H|2|O}} → 4 {{chem|Fe|3+}} + 20 {{chem|OH|-}} + 3 O<sub>2</sub>

The Fe<sup>3+</sup> ion has a large simple cationic chemistry, although the pale-violet hexaquo ion {{chem2|[Fe(H2O)6](3+)}} is very readily hydrolyzed when pH increases above 0 as follows:<ref name="Greenwood1088">Greenwood and Earnshaw, pp. 1088–91</ref>

{|
|-
| {{chem2|[Fe(H2O)6](3+)}} || {{eqm}} {{chem2|[Fe(H2O)5(OH)](2+) + H(+)}} || ''[[Chemistry: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 [[Chemistry:Iron(III) oxide|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 [[Chemistry: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.<ref name="Greenwood1088" /> 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 [[Chemistry:Carbonate|carbonate]] anions are added, which instead results in white [[Chemistry:Iron(II) carbonate|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 [[Chemistry:Iron(III) oxide|iron(III) oxide]] that accounts for the brown deposits present in a sizeable number of streams.<ref name="Greenwood1091">Greenwood and Earnshaw, pp. 1091–97</ref>

==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 [[Chemistry:Enantiomer|enantiomorphs]] of the ferrioxalate ion]]
Many coordination compounds of iron are known, although few exhibit [[Chemistry:High-valent iron|high-valent iron]].  A typical six-coordinate anion is hexachloroferrate(III), [FeCl<sub>6</sub>]<sup>3−</sup>, found in the mixed [[Chemistry:Salt|salt]] [[Chemistry:Tetrakis(methylammonium) hexachloroferrate(III) chloride|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 |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 isomers. 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}} [[Chemistry:Moiety|moiety]].<ref>{{cite journal<!--Deny Citation bot-->|last1 = Giannoccaro|first1 = P.|last2 = Sacco|first2 = A.| title=Bis&#91;Ethylenebis(Diphenylphosphine)&#93;‐Hydridoiron Complexes |journal = [[Chemistry:Inorganic Syntheses|Inorg. Synth.]]|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 [[Chemistry:Oxalate|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.<ref name="Greenwood1088" /> [[Chemistry:Potassium ferrioxalate|Potassium ferrioxalate]] is used in chemical actinometry and along with its [[Chemistry:Sodium ferrioxalate|sodium salt]] undergoes photoreduction applied in old-style photographic processes. The [[Chemistry:Hydrate|dihydrate]] of [[Chemistry:Iron(II) oxalate|iron(II) oxalate]] has a [[Chemistry:Polymer|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:Physics and Chemistry of Minerals|Phys. Chem. Minerals]]|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 [[Chemistry:Chromium|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 [[Chemistry:Phenol|phenol]]s or [[Biology:Enol|enol]]s. For example, in the [[Chemistry:Ferric chloride test|ferric chloride test]], used to determine the presence of phenols, [[Chemistry:Iron(III) chloride|iron(III) chloride]] reacts with a phenol to form a deep violet complex:<ref name="Greenwood1088" />

:3 ArOH + FeCl<sub>3</sub> → Fe(OAr)<sub>3</sub> + 3 HCl (Ar = [[Chemistry:Aryl|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). [[Chemistry:Thiocyanate|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 [[Chemistry:Spectrochemical series|spectrochemical series]] such as [[Chemistry:Cyanide|cyanide]]. An example of a low-spin iron(III) complex is [Fe(CN)<sub>6</sub>]<sup>3−</sup>. The cyanide ligands may easily be detached in [Fe(CN)<sub>6</sub>]<sup>3−</sup>, and hence this complex is poisonous, unlike the iron(II) complex [Fe(CN)<sub>6</sub>]<sup>4−</sup> found in Prussian blue,<ref name="Greenwood1088" /> which does not release [[Chemistry:Hydrogen cyanide|hydrogen cyanide]] except when dilute acids are added.<ref name="Greenwood1091" /> Iron shows a great variety of electronic [[Physics: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.<ref name="Greenwood1079" />

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.<ref name="Greenwood1091" />

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

[[File:Prussian blue.jpg|thumb|right|Prussian blue]]
[[Chemistry: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 [[Chemistry:Potassium ferricyanide|potassium ferricyanide]] and [[Chemistry:Potassium ferrocyanide|potassium ferrocyanide]] to form Prussian blue.<ref name="HollemanAF" />

Another old example of an organoiron compound is [[Chemistry:Iron pentacarbonyl|iron pentacarbonyl]], Fe(CO)<sub>5</sub>, in which a neutral iron atom is bound to the carbon atoms of five [[Chemistry:Carbon monoxide|carbon monoxide]] molecules. The compound can be used to make [[Chemistry:Carbonyl iron|carbonyl iron]] powder, a highly reactive form of metallic iron. [[Physics:Thermal decomposition|Thermolysis]] of iron pentacarbonyl gives [[Chemistry:Triiron dodecacarbonyl|triiron dodecacarbonyl]], {{chem2|Fe3(CO)12}}, a complex with a cluster of three iron atoms at its core. Collman's reagent, [[Chemistry:Disodium tetracarbonylferrate|disodium tetracarbonylferrate]], is a useful reagent for organic chemistry; it contains iron in the −2 oxidation state. [[Chemistry:Cyclopentadienyliron dicarbonyl dimer|Cyclopentadienyliron dicarbonyl dimer]] contains iron in the rare +1 oxidation state.<ref>{{Greenwood&Earnshaw1st|pages=1282–86}}.</ref>

{{multiple image
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A landmark in this field was the discovery in 1951 of the remarkably stable sandwich compound [[Chemistry:Ferrocene|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 |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= [[Organization: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 [[Biography:Robert Burns Woodward|Woodward]] and [[Biography:Geoffrey Wilkinson|Wilkinson]]<ref>{{cite journal |first1= G. |last1=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 = [[Organization:Journal of the American Chemical Society|J. Am. Chem. Soc.]] |year = 1952|volume = 74 |pages = 2125–2126 |doi = 10.1021/ja01128a527 |issue = 8}}</ref> and [[Biography: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.<ref name="Greenwood1104">Greenwood and Earnshaw, p. 1104</ref>

Iron-centered organometallic species are used as [[Chemistry:Catalyst|catalyst]]s. The [[Chemistry:Knölker complex|Knölker complex]], for example, is a [[Chemistry:Transfer hydrogenation|transfer hydrogenation]] catalyst for [[Chemistry:Ketone|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 [[Chemistry:Iron(II) sulfate|iron(II) sulfate]] (FeSO<sub>4</sub>·7[[Physics:Water of crystallization|H<sub>2</sub>O]]) and [[Chemistry:Iron(III) chloride|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" />

==See also==
* [[Chemistry:Cobalt compounds|Cobalt compounds]]
* [[Chemistry:Ferberite|Ferberite]]

==References==
{{Reflist}}

{{Iron compounds}}

[[Category:Iron]]
[[Category:Iron compounds]]
[[Category:Chemical compounds by element]]

{{Sourceattribution|Iron compounds}}