Ferric

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Ferric oxide, commonly, though not precisely, called rust

Ferric refers to iron-containing materials or compounds. In chemistry the term is reserved for iron with an oxidation number of +3, also denoted iron(III) or Fe3+. On the other hand, ferrous refers to iron with oxidation number of +2, denoted iron(II) or Fe2+. Iron(III) is usually the most stable form of iron in air, as illustrated by the pervasiveness of rust, an insoluble iron(III)-containing material.

Ferric iron and life[edit]

The bioavailability of iron is of great interest because (i) all known forms of life require iron and (ii) ordinary iron(III) compounds are insoluble in an aerobic environment. Iron-deficiency Anemia illustrates the problems resulting from low iron intake. Many foods contain soluble iron compounds and are therefore necessary for good nutrition.

The low bioavailability of iron affects all forms of life. Bacteria secrete iron-attracting agents called siderophores that form soluble compounds of iron that can be reabsorbed into the cell and used in building iron-containing metalloproteins. The impact of increasing the bioavailability of iron was famously demonstrated by an experiment where a large area of the ocean surface was sprayed with iron(III) salts. After several days, the phytoplankton within the treated area bloomed to such an extent that the effect was visible from outer space. This fertilizing process has been proposed as a means to mitigate the carbon dioxide content of the atmosphere.[1]

Ferric iron mitigates the eutrophication of lakes by reducing the bioavailability of phosphorus (as phosphate) in the water. Mitigation arises because ferric phosphate is insoluble. Like iron, phosphate is often a limiting nutrient, and its reduction in concentration from solution limits the growth of algae, which in turn prevents eutrophication.

Hydrolysis of iron(III) and rust[edit]

In water, ferric iron forms compounds that are often insoluble, at least near neutral pH. A salt of ferric iron hydrolyzes water and produces iron(III) oxide-hydroxides while contributing hydrogen ions to the solution, lowering the pH. In contrast, typical Na+ salts (e.g. NaCl) dissolve in water without lowering the pH.[2] The differing behavior of Na+ vs Fe3+ ions reflects the effect of charge: water when bound to Fe3+ is highly acidic, inducing the formation of iron(III) hydroxides, which polymerize via the process called olation. Aluminium(III) (Al3+) behaves similarly to ferric ion.

Rust, a mixture of ferric hydroxide compounds, illustrates the low solubility of ferric ions in water. Various reagents cause rust to dissolve even at neutral pH. These ligands include EDTA, which forms a chelate complex with the ion, displacing the hydroxide and oxide ligands that comprise rust. For this reason, EDTA is often used to dissolve iron deposits or to deliver soluble iron in fertilizers. Citrate also solubilizes ferric ion at neutral pH, although its complexes are less stable than those of EDTA.

Inorganic chemistry[edit]

Ferric iron is a d5 center, meaning that the metal has five "valence" electrons" in the 3d orbital shell. The magnetism of ferric compounds is mainly determined by these five electrons, and their behavior depends on the number and type of ligands attached to iron, as described by ligand field theory. Usually ferric ions are surrounded by six ligands arranged in octahedron. Sometimes three and sometimes as many as seven ligands are observed. A common ferric compound is ferric chloride (FeCl3).

Biochemistry[edit]

Most iron-containing proteins contain ferric ions, at least transiently. Well studied examples include iron-sulfur clusters, oxyhemoglobin, ferritin, and the cytochromes.[3]

Etymology[edit]

The word ferric is derived from the Latin word "ferrum" for iron.

See also[edit]

References[edit]

  1. ^ Boyd PW, Watson AJ, Law CS, et al. (October 2000). "A mesoscale phytoplankton bloom in the polar Southern Ocean stimulated by iron fertilization". Nature 407 (6805): 695–702. doi:10.1038/35037500. PMID 11048709. 
  2. ^ Earnshaw, A.; Greenwood, N. N. (1997). Chemistry of the elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0-7506-3365-4. 
  3. ^ Berg, Jeremy Mark; Lippard, Stephen J. (1994). Principles of bioinorganic chemistry. Sausalito, Calif: University Science Books. ISBN 0-935702-73-3.