View source for Ferric

{{Short description|Iron in its +3 oxidation state}}
[[File:Potassium trioxalatoferrate(III) large crystals.jpg|thumb|[[Potassium ferrioxalate]] contains the iron(III) complex {{chem2|[Fe(C2O4)3](3-)}}.]]

In [[chemistry]], '''iron(III)''' or '''''ferric''''' refers to the [[chemical element|element]] [[iron]] in its +3 [[oxidation number|oxidation state]].  ''Ferric chloride'' is an alternative name for [[iron(III) chloride]] ({{chem2|FeCl3}}). The adjective ''[[ferrous]]'' is used instead for [[iron(II)]] salts, containing the cation Fe<sup>2+</sup>. The word ''[[wikt:ferric|ferric]]'' is derived from the [[Latin]] word {{wikt-lang|la|ferrum}}, meaning "iron".  

Although often abbreviated as '''Fe<sup>3+</sup>''', that naked ion does not exist except under extreme conditions. Iron(III) centres are found in many compounds and [[coordination complex]]es, where Fe(III) is bonded to several [[Ligand|ligands]]. A molecular ferric complex is the [[anion]] [[ferrioxalate]], {{chem2|[Fe(C2O4)3](3-)}}, with three [[bidentate]] [[oxalate]] ions surrounding the Fe core.  Relative to lower oxidation states, ferric is less common in [[organoiron chemistry]], but the [[ferrocenium]] cation {{chem2|[Fe(C5H5)2](+)}} is well known.

==Iron(III) in biology==
All known forms of life require iron, which usually exists in Fe(II) or Fe(III) oxidation states.<ref>{{cite web | title = Iron integral to the development of life on Earth – and the possibility of life on other planets | url = https://www.ox.ac.uk/news/2021-12-07-iron-integral-development-life-earth-and-possibility-life-other-planets | date = 7 December 2021 | publisher = [[University of Oxford]] | accessdate = 9 May 2022}}</ref> Many [[protein]]s in living beings contain iron(III) centers.  Examples of such [[metalloprotein]]s include [[oxyhemoglobin]], [[ferredoxin]], and the [[cytochrome]]s.  Many organisms, from bacteria to humans, store iron as microscopic crystals (3 to 8 nm in diameter) of [[iron(III) oxide hydroxide]], inside  a shell of the protein [[ferritin]], from which it can be recovered as needed.<ref>{{cite book |author=Berg, Jeremy Mark |author2=Lippard, Stephen J. |title=Principles of bioinorganic chemistry |publisher=University Science Books |location=Sausalito, Calif |year=1994 |isbn=0-935702-73-3 }}</ref>

Insufficient iron in the human diet causes [[anemia]].  Animals and humans can obtain the necessary iron from foods that contain it in assimilable form, such as meat. Other organisms must obtain their iron from the environment.  However, iron tends to form highly insoluble iron(III) oxides/hydroxides in aerobic ([[oxygen]]ated) environment, especially in [[calcareous soil]]s. [[Bacteria]] and [[graminaceae|grass]]es can thrive in such environments by secreting compounds called [[siderophore]]s that form soluble complexes with iron(III), that can be reabsorbed into the cell. (The other plants instead encourage the growth around their roots of certain bacteria that [[redox|reduce]] iron(III) to the more soluble iron(II).)<ref name=marsch94>H. Marschner and V. Römheld (1994): "Strategies of plants for acquisition of iron". ''Plant and Soil'', volume 165, issue 2, pages 261–274. {{doi|10.1007/BF00008069}}</ref>  

The insolubility of iron(III) compounds is also responsible for the low levels of iron in seawater, which is often the limiting factor for the growth of the microscopic plants ([[phytoplankton]]) that are the basis of the marine food web.<ref>{{cite journal |vauthors=Boyd PW, Watson AJ, Law CS |title=A mesoscale phytoplankton bloom in the polar Southern Ocean stimulated by iron fertilization |journal=Nature |volume=407 |issue=6805 |pages=695–702 |date=October 2000 |pmid=11048709 |doi=10.1038/35037500 |display-authors=etal|bibcode=2000Natur.407..695B |s2cid=4368261 }}</ref>

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

==Iron(III) salts and complexes==
Typically iron(III) salts, like the "[[iron(III) chloride|chloride]]" are [[aquo complex]]es with the formulas {{chem2|[Fe(H2O)5Cl](2+), [Fe(H2O)4Cl2]+, and [Fe(H2O)3Cl3]}}.  [[Iron(III) nitrate]] and [[iron(III) perchlorate]] are thought to initially dissolve in water to give {{chem2|[Fe(H2O)6](3+)}} ions.  In these complexes, the protons are acidic.  Eventually these complexes [[hydrolysis|hydrolyze]] producing [[iron(III) hydroxide]]s {{chem2|Fe(OH)3 or Fe(O)OH}} that continue to react, in part via the process called [[olation]].  These hydroxides [[precipitate]] out of the solution or form colloids.  These reactions liberate [[hydrogen]] ions {{chem2|H(+)}} lowering the pH of its solutions. The [[chemical equilibrium|equilibria]] are elaborate:<ref name=earn>{{cite book |author=Earnshaw, A. |author2=Greenwood, N. N. |title=Chemistry of the elements |publisher=Butterworth-Heinemann |location=Oxford |year=1997 |isbn=0-7506-3365-4 |edition=2nd}}</ref><ref>{{cite journal |doi=10.1002/jctb.280460204 |title=Review of the hydrolysis of iron(III) and the crystallization of amorphous iron(III) hydroxide hydrate |date=1989 |last1=Cornell |first1=Rochelle M. |last2=Giovanoli |first2=Rudolf |last3=Schneider |first3=Walter |journal=Journal of Chemical Technology & Biotechnology |volume=46 |issue=2 |pages=115–134 |bibcode=1989JCTB...46..115C }}</ref><ref>{{cite journal |doi=10.1021/ic980813c |title=New Reaction Path in the Dissociation of the Fe<sub>2</sub>(μ-OH)<sub>2</sub>(H<sub>2</sub>O)<sub>8</sub><sup>4+</sup>Complex |date=1999 |last1=Lente |first1=Gábor |last2=Fábián |first2=István |journal=Inorganic Chemistry |volume=38 |issue=3 |pages=603–605 |pmid=11673970 }}</ref>
: {{chem2|[Fe(H2O)6](3+) <->  [Fe(H2O)5OH](2+)  +  H+ }}
: {{chem2|2 [Fe(H2O)5OH](2+)  <->  [Fe2(H2O)4(OH)2](4+)  +  2H2O }}
: {{chem2|2 [Fe(H2O)4(OH)2]+  <->  [Fe2(H2O)8(OH)2]2+  + 2 H2O }}

The aquo ligands on iron(III) complexes are labile.  This behavior is visualized by the color change brought about by reaction with thiocyanate to give a deep red [[thiocyanate complex]].

===Iron(III) with organic ligands===
In the presence [[chelation|chelating]] ligands, the complex hydrolysis reactions are avoided.  One of these ligands is [[EDTA]], which is often used to dissolve iron deposits or added to fertilizers to make iron in the soil available (soluble) to plants. [[Citrate]] also solubilizes ferric ion at neutral pH, although its complexes are less stable than those of EDTA.  Many chelating ligands - the [[siderophore]]s - are produced naturally to dissolve iron(III) oxides.

Iron(III) complexes with [[2,2'-Bipyridine|1,10-phenanthrolinebipyridine]] is soluble and can sustain reduction to it iron(II) derivative:
[[File:Fe(bipy)3 redox.svg|thumb|360px|center|Redox reaction of [Fe(bipyridine)<sub>3</sub>]<sup>3+</sup>.]]

==Iron(III) minerals and other solids==
[[File:Iron(III)-oxide-sample.jpg|thumb|[[Ferric oxide]], commonly called [[rust]], is a very complicated material that contains iron(III).]]

Iron(III) is found in many minerals and solids, e.g., [[iron(III) oxide|oxide]]  {{chem2|Fe2O3}} (hematite) and [[iron(III) oxide-hydroxide]] {{chem2|FeO(OH)}} are extremely insoluble reflecting their [[polymer|polymeric]] structure. [[Rust]] is a mixture of iron(III) oxide and oxide-hydroxide that usually forms when iron metal is exposed to [[humidity|humid]] air.  Unlike the [[Passivation (chemistry)|passivating]] oxide layers that are formed by other metals, like [[chromium]] and [[aluminum]], rust flakes off, because it is bulkier than the metal that formed it.  Therefore, unprotected iron objects will in time be completely turned into rust.

==Bonding==
[[File:L.s. vs h.s. d5 octahedral.svg|thumb|left|d-orbital splitting scheme for low- and high spin octahedral Fe(III) complex]]
Iron(III) is a d<sup>5</sup> center, meaning that the metal has five "valence" electrons in the 3d orbital shell. The number and type of ligands bound to iron(III) determine how these electrons arrange themselves.  With so-called "strong field ligands" such as [[cyanide]], the five electrons pair up as best they can.  Thus [[ferricyanide]] ({{chem2|[Fe(CN)6](3-)}} has only one unpaired electron.  It is low-spin. With so-called "weak field ligands" such as [[water]], the five electrons are unpaired.  Thus [[aquo complex]] ({{chem2|[Fe(H2O)6](3+)}} has only five unpaired electrons.  It is high-spin.  With chloride, iron(III) forms tetrahedral complexes, e.g.  ({{chem2|[Fe(Cl)4]-}}.  Tetrahedral complexes are high spin. The magnetism of ferric complexes can show when they are high or low spin.

==See also==

* {{annotated link|Ferric chloride}} ([[Iron(III) chloride]])
* {{annotated link|Ferric oxide}} ([[Iron(III) oxide]])
* {{annotated link|Ferric fluoride}} ([[Iron(III) fluoride]])
* {{annotated link|Ferrous}}

==References==

{{reflist}}

[[Category:Iron(III) compounds]]