In
chemistry, iron(II) refers to the
elementiron in its +2
oxidation state. The adjective ferrous or the prefix ferro- is often used to specify such compounds, as in ferrous chloride for
iron(II) chloride (FeCl2). The adjective ferric is used instead for
iron(III) salts, containing the cation Fe3+. The word ferrous is derived from the
Latin word ferrum, meaning "iron".
In
ionic compounds (salts), such an atom may occur as a separate
cation (positive ion) abbreviated as Fe2+, although more precise descriptions include other ligands such as water and halides. Iron(III) centres occur in
coordination complexes, such as in the
anionferrocyanide, [Fe(CN)64−, where six
cyanide ligands are bound the metal centre; or, in
organometallic compounds, such as the
ferrocene[Fe(C2H5)2, where two
cyclopentadienyl anions are bound to the FeII centre.
All known forms of life require iron.[1] Many
proteins in living beings contain iron(III) centers. Examples of such
metalloproteins include
hemoglobin,
ferredoxin, and the
cytochromes. In many of these proteins, Fe(II) converts reversibly to Fe(III) [2]
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 (
oxygenated) environment, especially in
calcareous soils.
Bacteria and
grasses can thrive in such environments by secreting compounds called
siderophores 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
reduce 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]
The aquo ligands on iron(II) complexes are labile. It reacts with
1,10-phenanthroline to give the blue iron(II) derivative:
When metallic iron (oxidation state 0) is placed in a solution of
hydrochloric acid, iron(II) chloride is formed, with release of
hydrogen gas, by the reaction
Fe0 + 2 H+ → Fe2+ + H2
Iron(II) is oxidized by hydrogen peroxide to
iron(III), forming a
hydroxyl radical and a
hydroxide ion in the process. This is the
Fenton reaction. Iron(III) is then reduced back to iron(II) by another molecule of hydrogen peroxide, forming a
hydroperoxyl radical and a
proton. The net effect is a
disproportionation 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)
The
free radicals generated by this process engage in secondary reactions, which can degrade many organic and biochemical compounds.
Iron(II) minerals and other solids
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 often
nonstoichiometric. For example, "ferrous sulfide" can refer to the 1:1 species (mineral name
troilite) or a host of Fe-deficient derivatives (
pyrrhotite). The mineral
magnetite ("lode stone") is a mixed-valence compound with both Fe(II) and Fe(III), Fe3O4.
Bonding
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 so-called "strong field ligands" such as
cyanide, the six electrons pair up. Thus
ferrocyanide ([Fe(CN)64− has no unpaired electrons. It is low-spin. With so-called "weak field ligands" such as
water, the four of the six electrons are unpaired. Thus
aquo complex ([Fe(H2O)62+ is paramagnetic. It is high-spin. With chloride, iron(III) forms tetrahedral complexes, e.g. [FeCl42−. Tetrahedral complexes are high spin.
Ferric – The element iron in its +3 oxidation state — [ Iron(III)] compounds
Ferromagnetism – Mechanism by which materials form into and are attracted to magnets
Ferrous metal recycling – Recyclable materials left over from manufactured products after their usePages displaying short descriptions of redirect targets
Iron(II) oxide – Inorganic compound with the formula FeO (ferrous oxide)
Iron(II) bromide – chemical compoundPages displaying wikidata descriptions as a fallback (ferrous bromide)
Steelmaking – Process for producing steel from iron ore and scrap
^Berg, Jeremy Mark; Lippard, Stephen J. (1994). Principles of bioinorganic chemistry. Sausalito, Calif: University Science Books.
ISBN0-935702-73-3.
^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
^Petsch, S.T. (2014). "10.11 - The Global Oxygen Cycle". In Holland, H.D.; Turekian, K.K. (eds.). Treatise on Geochemistry. Reference Module in Earth Systems and Environmental Sciences. Vol. 10 (Second ed.). Elsevier. pp. 437–473.
doi:
10.1016/B978-0-08-095975-7.00811-1.
ISBN978-0-08-095975-7.
^Earnshaw, A.; Greenwood, N. N. (1997). Chemistry of the elements (2nd ed.). Oxford: Butterworth-Heinemann.
ISBN0-7506-3365-4.