Ferrous

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Iron(II) chloride tetrahydrate, .mw-parser-output .template-chem2-su{display:inline-block;font-size:80%;line-height:1;vertical-align:-0.35em}.mw-parser-output .template-chem2-su>span{display:block;text-align:left}.mw-parser-output sub.template-chem2-sub{font-size:80%;vertical-align:-0.35em}.mw-parser-output sup.template-chem2-sup{font-size:80%;vertical-align:0.65em} FeCl2*4H2O. Eisen(II)-chlorid-Tetrahydrat.jpg
Iron(II) chloride tetrahydrate, FeCl2·4H2O.

In chemistry, iron(II) refers to the element iron 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".

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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(II) centres occur in coordination complexes, such as in the anion ferrocyanide, [Fe(CN)6]4−, 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.

Ferrous ions in biology

All known forms of life require iron. [1] Many proteins in living beings contain iron(II) 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]

Pourbaix diagram of aqueous iron Pourbaix Diagram of Iron.svg
Pourbaix diagram of aqueous iron

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]

Ferrous salts and complexes

Typically iron(II) salts, like the "chloride" are aquo complexes with the formulas [Fe(H2O)6]2+, as found in ferrous ammonium sulfate. [5]

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]

The free radicals generated by this process engage in secondary reactions, which can degrade many organic and biochemical compounds.

Redox reaction of [Fe(bipyridine)3] . Fe(bipy)3 redox.svg
Redox reaction of [Fe(bipyridine)3] .

Ferrous minerals and other solids

Iron(II) oxide (ferrous oxide), FeO, is a very complicated material that contains iron(II). Iron(II) oxide.jpg
Iron(II) oxide (ferrous oxide), FeO, is a very complicated material that contains iron(II).

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

d-orbital splitting scheme for low- and high spin octahedral Fe(II) complexes. L.s. vs h.s. d6 octahedral.svg
d-orbital splitting scheme for low- and high spin octahedral Fe(II) complexes.

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 as cyanide, the six electrons pair up. Thus ferrocyanide ([Fe(CN)6]4− has no unpaired electrons, meaning it is a low-spin complex. With so-called "weak field ligands" such as water, four of the six electrons are unpaired, meaning it is a high-spin complex. Thus aquo complex [Fe(H2O)6]2+ is paramagnetic. With chloride, iron(II) forms tetrahedral complexes, e.g. [FeCl4]2−. Tetrahedral complexes are high-spin complexes.

See also

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<span class="mw-page-title-main">Ferric</span> The element iron in its +3 oxidation state

In chemistry, iron(III) or ferric refers to the element iron in its +3 oxidation state. Ferric chloride is an alternative name for iron(III) chloride (FeCl3). The adjective ferrous is used instead for iron(II) salts, containing the cation Fe2+. The word ferric is derived from the Latin word ferrum, meaning "iron".

<span class="mw-page-title-main">Prussian blue</span> Synthetic pigment

Prussian blue (also known as Berlin blue, Brandenburg blue, Parisian and Paris blue) is a dark blue pigment produced by oxidation of ferrous ferrocyanide salts. It has the chemical formula Fe3+4[Fe2+(CN)6]3. Turnbull's blue is essentially identical chemically, excepting that it has different impurities and particle sizes—because it is made from different reagents—and thus it has a slightly different color.

<span class="mw-page-title-main">Potassium ferricyanide</span> Chemical compound

Potassium ferricyanide is the chemical compound with the formula K3[Fe(CN)6]. This bright red salt contains the octahedrally coordinated [Fe(CN)6]3− ion. It is soluble in water and its solution shows some green-yellow fluorescence. It was discovered in 1822 by Leopold Gmelin.

In chemistry, a reducing agent is a chemical species that "donates" an electron to an electron recipient.

Iron(III) chloride describes the inorganic compounds with the formula FeCl3(H2O)x. Also called ferric chloride, these compounds are some of the most important and commonplace compounds of iron. They are available both in anhydrous and in hydrated forms, which are both hygroscopic. They feature iron in its +3 oxidation state. The anhydrous derivative is a Lewis acid, while all forms are mild oxidizing agents. It is used as a water cleaner and as an etchant for metals.

Iron(II) chloride, also known as ferrous chloride, is the chemical compound of formula FeCl2. It is a paramagnetic solid with a high melting point. The compound is white, but typical samples are often off-white. FeCl2 crystallizes from water as the greenish tetrahydrate, which is the form that is most commonly encountered in commerce and the laboratory. There is also a dihydrate. The compound is highly soluble in water, giving pale green solutions.

Iron–sulfur proteins are proteins characterized by the presence of iron–sulfur clusters containing sulfide-linked di-, tri-, and tetrairon centers in variable oxidation states. Iron–sulfur clusters are found in a variety of metalloproteins, such as the ferredoxins, as well as NADH dehydrogenase, hydrogenases, coenzyme Q – cytochrome c reductase, succinate – coenzyme Q reductase and nitrogenase. Iron–sulfur clusters are best known for their role in the oxidation-reduction reactions of electron transport in mitochondria and chloroplasts. Both Complex I and Complex II of oxidative phosphorylation have multiple Fe–S clusters. They have many other functions including catalysis as illustrated by aconitase, generation of radicals as illustrated by SAM-dependent enzymes, and as sulfur donors in the biosynthesis of lipoic acid and biotin. Additionally, some Fe–S proteins regulate gene expression. Fe–S proteins are vulnerable to attack by biogenic nitric oxide, forming dinitrosyl iron complexes. In most Fe–S proteins, the terminal ligands on Fe are thiolate, but exceptions exist.

<span class="mw-page-title-main">Iron(II) hydroxide</span> Chemical compound

Iron (II) hydroxide or ferrous hydroxide is an inorganic compound with the formula Fe(OH)2. It is produced when iron (II) salts, from a compound such as iron(II) sulfate, are treated with hydroxide ions. Iron(II) hydroxide is a white solid, but even traces of oxygen impart a greenish tinge. The air-oxidised solid is sometimes known as "green rust".

<span class="mw-page-title-main">Ferricyanide</span> Anion in which a Fe3+ ion is complexed by 6 CN− ions

Ferricyanide is the name of the anion [Fe(CN)6]3−. It is also called hexacyanoferrate(III) and in rare, but systematic nomenclature, hexacyanidoferrate(III). The most common salt of this anion is potassium ferricyanide, a red crystalline material that is used as an oxidant in organic chemistry.

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<span class="mw-page-title-main">Iron(III) nitrate</span> Chemical compound

Iron(III) nitrate, or ferric nitrate, is the name used for a series of inorganic compounds with the formula Fe(NO3)3.(H2O)n. Most common is the nonahydrate Fe(NO3)3.(H2O)9. The hydrates are all pale colored, water-soluble paramagnetic salts.

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The Haber–Weiss reaction generates •OH (hydroxyl radicals) from H2O2 (hydrogen peroxide) and superoxide (•O2−) catalyzed by iron ions. It was first proposed by Fritz Haber and his student Joseph Joshua Weiss in 1932.

<span class="mw-page-title-main">Iron(III) acetate</span> Chemical compound

Ferric acetate is the iron compound with the formula Fe3O(O2CCH3)6(H2O)3]O2CCH3. This red brown solid is the acetate salt of the coordination complex [Fe3O(OAc)6(H2O)3]+ (OAc− is CH3CO2−). Commonly, the salt is known as "basic iron acetate". The formation of the red-brown complex was once used as a test for ferric ions.

<span class="mw-page-title-main">Sodium ferrocyanide</span> Chemical compound

Sodium ferrocyanide is the sodium salt of the coordination compound of formula [Fe(CN)6]4−. In its hydrous form, Na4Fe(CN)6·H2O, it is sometimes known as yellow prussiate of soda. It is a yellow crystalline solid that is soluble in water and insoluble in alcohol. The yellow color is the color of ferrocyanide anion. Despite the presence of the cyanide ligands, sodium ferrocyanide has low toxicity. The ferrocyanides are less toxic than many salts of cyanide, because they tend not to release free cyanide. However, like all ferrocyanide salt solutions, addition of an acid or exposure to UV light can result in the production of hydrogen cyanide gas, which is extremely toxic.

<span class="mw-page-title-main">Iron(III) sulfate</span> Chemical compound

Iron(III) sulfate (or ferric sulfate), is a family of inorganic compounds with the formula Fe2(SO4)3(H2O)n. A variety of hydrates are known, including the most commonly encountered form of "ferric sulfate". Solutions are used in dyeing as a mordant, and as a coagulant for industrial wastes. Solutions of ferric sulfate are also used in the processing of aluminum and steel.

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Cerimetry or cerimetric titration, also known as cerate oximetry, is a method of volumetric chemical analysis developed by Ion Atanasiu. It is a redox titration in which an iron(II)–1,10-phenanthroline complex (ferroin) color change indicates the end point. Ferroin can be reversibly discolored in its oxidized form upon titration with a Ce4+ solution. The use of cerium(IV) salts as reagents for volumetric analysis was first proposed in the middle of 19th century, but systematic studies did not start until about 70 years later. Standard solutions can be prepared from different Ce4+ salts, but often cerium sulfate is chosen.

<span class="mw-page-title-main">Iron(II) perchlorate</span> Chemical compound

Iron(II) perchlorate is the inorganic compound with the formula Fe(ClO4)2·6H2O. A green, water-soluble solid, it is produced by the reaction of iron metal with dilute perchloric acid followed by evaporation of the solution:

References

  1. ↑ "Iron integral to the development of life on Earth – and the possibility of life on other planets". University of Oxford. 7 December 2021. Retrieved 9 May 2022.
  2. ↑ Berg, Jeremy Mark; Lippard, Stephen J. (1994). Principles of bioinorganic chemistry. Sausalito, Calif: University Science Books. ISBN   0-935702-73-3.
  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
  4. ↑ 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. ISBN   978-0-08-095975-7.
  5. ↑ Earnshaw, A.; Greenwood, N. N. (1997). Chemistry of the elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN   0-7506-3365-4.
  6. ↑ Tang, Zhongmin; Zhao, Peiran; Wang, Han; Liu, Yanyan; Bu, Wenbo (2021). "Biomedicine Meets Fenton Chemistry". Chemical Reviews. 121 (4): 1981–2019. doi:10.1021/acs.chemrev.0c00977. PMID   33492935. S2CID   231712587.
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