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Inchemistry,iron(II) refers to theelementiron in its +2oxidation state. The adjectiveferrous or the prefixferro- is often used to specify such compounds, as inferrous chloride foriron(II) chloride (FeCl2). The adjectiveferric is used instead foriron(III) salts, containing the cation Fe3+. The wordferrous is derived from theLatin wordferrum, meaning "iron".
Inionic compounds (salts), such an atom may occur as a separatecation (positive ion) abbreviated asFe2+, although more precise descriptions include other ligands such as water and halides. Iron(II) centres occur incoordination complexes, such as in theanionferrocyanide,[Fe(CN)6]4−, where sixcyanide ligands are bound the metal centre; or, inorganometallic compounds, such as theferrocene[Fe(C2H5)2], where twocyclopentadienyl anions are bound to the FeII centre.
All known forms of life require iron.[1] Manyproteins in living beings contain iron(II) centers. Examples of suchmetalloproteins includehemoglobin,ferredoxin, and thecytochromes. In many of these proteins, Fe(II) converts reversibly to Fe(III).[2]
Insufficient iron in the human diet causesanemia. 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 (oxygenated) environment, especially incalcareous soils.Bacteria andgrasses can thrive in such environments by secreting compounds calledsiderophores 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 thatreduce iron(III) to the more soluble iron(II).)[3]
In contrast to iron(III) aquo complexes, iron(II) aquo complexes are soluble in water near neutral pH.[citation needed] Ferrous iron is, however, oxidized by the oxygen in air, converting to iron(III).[4]
Typically iron(II) salts, like the "chloride" areaquo complexes with the formulas[Fe(H2O)6]2+, as found inMohr's salt.[5]
The aquo ligands on iron(II) complexes are labile. It reacts with1,10-phenanthroline to give the blue iron(II) derivative:
When metallic iron (oxidation state 0) is placed in a solution ofhydrochloric acid, iron(II) chloride is formed, with release ofhydrogen gas, by the reaction
Iron(II) is oxidized by hydrogen peroxide toiron(III), forming ahydroxyl radical and ahydroxide ion in the process. This is theFenton reaction. Iron(III) is then reduced back to iron(II) by another molecule of hydrogen peroxide, forming ahydroperoxyl radical and aproton. The net effect is adisproportionation of hydrogen peroxide to create two different oxygen-radical species, with water (H+ + OH−) as a byproduct.[6]
| Fe2+ + H2O2 → Fe3+ + HO• + OH− | 1 |
| Fe3+ + H2O2 → Fe2+ + HOO• + H+ | 2 |
Thefree radicals generated by this process engage in secondary reactions, which can degrade many organic and biochemical compounds.
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Iron(II) is found in many minerals and solids. Examples include the sulfide and oxide, FeS and FeO. These formulas are deceptively simple because these sulfides and oxides are oftennonstoichiometric. For example, "ferrous sulfide" can refer to the 1:1 species (mineral nametroilite) or a host of Fe-deficient derivatives (pyrrhotite). The mineralmagnetite ("lode stone") is a mixed-valence compound with both Fe(II) and Fe(III), Fe3O4.
Iron(II) is a d6 center, meaning that the metal has six "valence" electrons in the 3d orbital shell. The number and type of ligands bound to iron(II) determine how these electrons arrange themselves. With the so-called "strong field ligands" such ascyanide, the six electrons pair up. Thusferrocyanide ([Fe(CN)6]4− has no unpaired electrons, meaning it is a low-spin complex. With so-called "weak field ligands" such aswater, four of the six electrons are unpaired, meaning it is ahigh-spin complex. Thusaquo complex[Fe(H2O)6]2+ isparamagnetic. Withchloride, iron(II) forms tetrahedral complexes, e.g.[FeCl4]2−. Tetrahedral complexes are high-spin complexes.
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