Editing Iron (section)

===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}}