Ferric
In chemistry, iron or ferric refers to the element iron in its +3 oxidation state. Ferric chloride is an alternative name for iron chloride. The adjective ferrous is used instead for iron salts, containing the cation Fe2+. The word ferric is derived from the Latin word, meaning "iron".
Although often abbreviated as Fe3+, that naked ion does not exist except under extreme conditions. Iron centres are found in many compounds and coordination complexes, where Fe is bonded to several ligands. A molecular ferric complex is the anion ferrioxalate,, 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 is well known.
Iron(III) in biology
All known forms of life require iron, which usually exists in Fe or Fe oxidation states. Many proteins in living beings contain iron centers. Examples of such metalloproteins include oxyhemoglobin, ferredoxin, and the cytochromes. Many organisms, from bacteria to humans, store iron as microscopic crystals of iron oxide hydroxide, inside a shell of the protein ferritin, from which it can be recovered as needed.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 oxides/hydroxides in aerobic environment, especially in calcareous soils. Bacteria and grasses can thrive in such environments by secreting compounds called siderophores that form soluble complexes with iron, that can be reabsorbed into the cell. to the more soluble iron
The insolubility of iron 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 that are the basis of the marine food web.
Iron(III) salts and complexes
Typically iron salts, like the "chloride" are aquo complexes with the formulas. Iron nitrate and iron perchlorate are thought to initially dissolve in water to give ions. In these complexes, the protons are acidic. Eventually these complexes hydrolyze producing iron hydroxides 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 lowering the pH of its solutions. The equilibria are elaborate:The aquo ligands on iron 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 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 to plants. Citrate also solubilizes ferric ion at neutral pH, although its complexes are less stable than those of EDTA. Many chelating ligands - the siderophores - are produced naturally to dissolve iron oxides.Iron complexes with 1,10-phenanthrolinebipyridine is soluble and can sustain reduction to it iron derivative:
File:Fe3 redox.svg|thumb|360px|center|Redox reaction of 3+.