Iron compounds

Chemical compounds containing iron
Oxidation state Representative compound −4 (d10s2) [FeIn6−xSnx] −2 (d10) Disodium tetracarbonylferrate (Collman's reagent) −1 (d9) 0 (d8) Iron pentacarbonyl 1 (d7) Cyclopentadienyliron dicarbonyl dimer ("Fp2") 2 (d6) Ferrous sulfate, ferrocene 3 (d5) Ferric chloride, ferrocenium tetrafluoroborate 4 (d4) , iron tetrafluoride 5 (d3) FeO43 6 (d2) Potassium ferrate7 (d1)[FeO4]– (matrix isolation, 4K)

Iron shows the characteristic chemical properties of the transition metals, namely the ability to form variable oxidation states differing by steps of one and a very large coordination and organometallic chemistry: indeed, it was the discovery of an iron compound, ferrocene, that revolutionalized the latter field in the 1950s. Iron is sometimes considered as a prototype for the entire block of transition metals, due to its abundance and the immense role it has played in the technological progress of humanity. Iron forms compounds mainly in the oxidation states +2 (iron(II), "ferrous") and +3 (iron(III), "ferric"). Iron also occurs in higher oxidation states, e. g. the purple potassium ferrate (K2FeO4), which contains iron in its +6 oxidation state. Although iron(VIII) oxide (FeO4) has been claimed, the report could not be reproduced and such a species from the removal of all electrons of the element beyond the preceding inert gas configuration (at least with iron in its +8 oxidation state) has been found to be improbable computationally. However, one form of anionic [FeO4]– with iron in its +7 oxidation state, along with an iron(V) peroxo isomer, has been detected by infrared spectroscopy at 4 K after cocondensation of laser ablated Fe atoms with a mixture of O2/Ar. Iron(IV) is a common intermediate in many biochemical oxidation reactions. Numerous organoiron compounds contain formal oxidation states of +1, 0, −1, or even −2. The oxidation states and other bonding properties are often assessed using the technique of Mössbauer spectroscopy. Many mixed valence compounds contain both iron(II) and iron(III) centers, such as magnetite and Prussian blue (|Fe4(Fe[CN]6)3). Iron is the first of the transition metals that cannot reach its group oxidation state of +8, although its heavier congeners ruthenium and osmium can, with ruthenium having more difficulty than osmium. Unlike many other metals, iron does not form amalgams with mercury. As a result, mercury is traded in standardized 76 pound flasks (34 kg) made of iron. Iron is by far the most reactive element in its group; it is pyrophoric when finely divided and dissolves easily in dilute acids, giving Fe2+. However, it does not react with concentrated nitric acid and other oxidizing acids due to the formation of an impervious oxide layer, which can nevertheless react with hydrochloric acid.

The binary ferrous and ferric halides are well known. The ferrous halides typically arise from treating iron metal with the corresponding hydrohalic acid to give the corresponding hydrated salts.

Due to its electronic structure, iron has a very large coordination and organometallic chemistry. Many coordination compounds of iron are known, although few exhibit high valent iron. A typical six coordinate anion is hexachloroferrate(III), [FeCl6]3−, found in the mixed salt tetrakis(methylammonium) hexachloroferrate(III) chloride. Complexes with multiple bidentate ligands have geometric isomers. For example, the trans chlorohydridobis(bis 1,2 (diphenylphosphino)ethane)iron(II) complex is used as a starting material for compounds with the |Fe(dppe)2 moiety. The ferrioxalate ion with three oxalate ligands (shown at right) displays helical chirality with its two non superposable geometries labelled Λ (lambda) for the left handed screw axis and Δ (delta) for the right handed screw axis, in line with IUPAC conventions. Iron(III) complexes are quite similar to those of chromium(III) with the exception of iron(III)'s preference for O donor instead of N donor ligands. The latter tend to be rather more unstable than iron(II) complexes and often dissociate in water. Many Fe–O complexes show intense colors and are used as tests for phenols or enols. For example, in the ferric chloride test, used to determine the presence of phenols, iron(III) chloride reacts with a phenol to form a deep violet complex:

A landmark in this field was the discovery in 1951 of the remarkably stable sandwich compound ferrocene |Fe(C5H5)2, by Pauson and Kealy and independently by Miller and colleagues, whose surprising molecular structure was determined only a year later by Woodward and Wilkinson and Fischer. Ferrocene is still one of the most important tools and models in this class. Iron centered organometallic species are used as catalysts. The Knölker complex, for example, is a transfer hydrogenation catalyst for ketones.

The iron compounds produced on the largest scale in industry are iron(II) sulfate (FeSO4·7H2O) and iron(III) chloride (FeCl3). The former is one of the most readily available sources of iron(II), but is less stable to aerial oxidation than Mohr's salt (|(NH4)2Fe(SO4)2*6H2O). Iron(II) compounds tend to be oxidized to iron(III) compounds in the air.