| Oxidation state |
Representative compound |
|---|---|
| −2 (d10) | Disodium tetracarbonylferrate (Collman's reagent) |
| −1 (d9) | Fe 2(CO)2− 8 |
| 0 (d8) | Iron pentacarbonyl |
| 1 (d7) | Cyclopentadienyliron dicarbonyl dimer ("Fp2") |
| 2 (d6) | Ferrous sulfate, ferrocene |
| 3 (d5) | Ferric chloride, ferrocenium tetrafluoroborate |
| 4 (d4) | Fe(diars) 2Cl2+ 2, iron tetrafluoride |
| 5 (d3) | FeO3− 4 |
| 6 (d2) | Potassium ferrate |
| 7 (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.[1] 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.[2] Its 26 electrons are arranged in the configuration [Ar]3d64s2, of which the 3d and 4s electrons are relatively close in energy, and thus it can lose a variable number of electrons and there is no clear point where further ionization becomes unprofitable.[3]
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.[4] 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.[5] Iron(IV) is a common intermediate in many biochemical oxidation reactions.[6][7] 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.[8] Many mixed valence compounds contain both iron(II) and iron(III) centers, such as magnetite and Prussian blue (Fe4(Fe[CN]6)3).[7] The latter is used as the traditional "blue" in blueprints.[9]
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.[10] Ruthenium exhibits an aqueous cationic chemistry in its low oxidation states similar to that of iron, but osmium does not, favoring high oxidation states in which it forms anionic complexes.[10] In the second half of the 3d transition series, vertical similarities down the groups compete with the horizontal similarities of iron with its neighbors cobalt and nickel in the periodic table, which are also ferromagnetic at room temperature and share similar chemistry. As such, iron, cobalt, and nickel are sometimes grouped together as the iron triad.[2]
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.[11]
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.[10] High purity iron, called electrolytic iron, is considered to be resistant to rust, due to its oxide layer.
- ^ Greenwood and Earnshaw, p. 905
- ^ a b Greenwood and Earnshaw, p. 1070
- ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. (p. 1074).
- ^ Huang, Wei; Xu, Wen-Hua; Schwarz, W.H.E.; Li, Jun (2016-05-02). "On the Highest Oxidation States of Metal Elements in MO4 Molecules (M = Fe, Ru, Os, Hs, Sm, and Pu)". Inorganic Chemistry. 55 (9): 4616–25. doi:10.1021/acs.inorgchem.6b00442. PMID 27074099.
- ^ Lu, Jun-Bo; Jian, Jiwen; Huang, Wei; Lin, Hailu; Li, Jun; Zhou, Mingfei (2016-11-16). "Experimental and theoretical identification of the Fe(VII) oxidation state in FeO4−". Phys. Chem. Chem. Phys. 18 (45): 31125–31131. Bibcode:2016PCCP...1831125L. doi:10.1039/c6cp06753k. PMID 27812577.
- ^ Nam, Wonwoo (2007). "High-Valent Iron(IV)–Oxo Complexes of Heme and Non-Heme Ligands in Oxygenation Reactions" (PDF). Accounts of Chemical Research. 40 (7): 522–531. doi:10.1021/ar700027f. PMID 17469792. Archived from the original (PDF) on 15 June 2021. Retrieved 22 February 2022.
- ^ a b Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils (1985). "Iron". Lehrbuch der Anorganischen Chemie (in German) (91–100 ed.). Walter de Gruyter. pp. 1125–46. ISBN 3-11-007511-3.
- ^ Reiff, William Michael; Long, Gary J. (1984). "Mössbauer Spectroscopy and the Coordination Chemistry of Iron". Mössbauer spectroscopy applied to inorganic chemistry. Springer. pp. 245–83. ISBN 978-0-306-41647-7.
- ^ Ware, Mike (1999). "An introduction in monochrome". Cyanotype: the history, science and art of photographic printing in Prussian blue. NMSI Trading Ltd. pp. 11–19. ISBN 978-1-900747-07-3.
- ^ a b c Cite error: The named reference
Greenwood1075was invoked but never defined (see the help page). - ^ Gmelin, Leopold (1852). "Mercury and Iron". Hand-book of chemistry. Vol. 6. Cavendish Society. pp. 128–29.