| 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 thetransition 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 theconfiguration [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 theoxidation states +2 (iron(II), "ferrous") and +3 (iron(III), "ferric"). Iron also occurs inhigher oxidation states, e.g. the purplepotassium 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] Numerousorganoiron 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 ofMössbauer spectroscopy.[8] Manymixed valence compounds contain both iron(II) and iron(III) centers, such asmagnetite andPrussian blue (Fe4(Fe[CN]6)3).[7] The latter is used as the traditional "blue" inblueprints.[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 neighborscobalt andnickel in the periodic table, which are also ferromagnetic atroom temperature and share similar chemistry. As such, iron, cobalt, and nickel are sometimes grouped together as theiron triad.[2]
Unlike many other metals, iron does not form amalgams withmercury. 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 ispyrophoric when finely divided and dissolves easily in dilute acids, giving Fe2+. However, it does not react with concentratednitric acid and other oxidizing acids due to the formation of an impervious oxide layer, which can nevertheless react withhydrochloric acid.[10] High purity iron, calledelectrolytic iron, is considered to be resistant to rust, due to its oxide layer.
Iron forms variousoxide and hydroxide compounds; the most common areiron(II,III) oxide (Fe3O4), andiron(III) oxide (Fe2O3).Iron(II) oxide also exists, though it is unstable at room temperature. Despite their names, they are actually allnon-stoichiometric compounds whose compositions may vary.[12] These oxides are the principal ores for the production of iron (seebloomery and blast furnace). They are also used in the production offerrites, usefulmagnetic storage media in computers, and pigments.
_oxide.jpg)
-oxide-sample.jpg)
The best known sulfide ispyrite (FeS2), also known as fool's gold owing to its golden luster.[7] It is not an iron(IV) compound, but is actually an iron(II)polysulfide containing Fe2+ andS2−
2 ions in a distortedsodium chloride structure.[12]
_chloride_hexahydrate.jpg)
The binary ferrous and ferrichalides are well-known. The ferrous halides typically arise from treating iron metal with the correspondinghydrohalic acid to give the corresponding hydrated salts.[7]
Iron reacts with fluorine, chlorine, and bromine to give the corresponding ferric halides,ferric chloride being the most common.[13]
Ferric iodide is an exception, being thermodynamically unstable due to the oxidizing power of Fe3+ and the high reducing power of I−:[13]
Ferric iodide, a black solid, is not stable in ordinary conditions, but can be prepared through the reaction ofiron pentacarbonyl withiodine andcarbon monoxide in the presence ofhexane and light at the temperature of −20 °C, with oxygen and water excluded.[13] Complexes of ferric iodide with some soft bases are known to be stable compounds.[14][15]
Thestandard reduction potentials in acidic aqueous solution for some common iron ions are given below:[10]
| Fe2+ + 2 e− | ⇌ Fe | E0 = −0.447 V |
| Fe3+ + 3 e− | ⇌ Fe | E0 = −0.037 V |
| FeO2− 4 + 8 H+ + 3 e− | ⇌ Fe3+ + 4 H2O | E0 = +2.20 V |
The red-purple tetrahedralferrate(VI) anion is such a strong oxidizing agent that it oxidizes nitrogen and ammonia at room temperature, and even water itself in acidic or neutral solutions:[13]
The Fe3+ ion has a large simple cationic chemistry, although the pale-violet hexaquo ion[Fe(H2O)6]3+ is very readily hydrolyzed when pH increases above 0 as follows:[16]
| [Fe(H2O)6]3+ | ⇌[Fe(H2O)5(OH)]2+ + H+ | K = 10−3.05 mol dm−3 |
| [Fe(H2O)5(OH)]2+ | ⇌[Fe(H2O)4(OH)2]+ + H+ | K = 10−3.26 mol dm−3 |
| 2[Fe(H2O)6]3+ | ⇌[Fe(H2O)4(OH)]4+2 + 2H+ + 2H2O | K = 10−2.91 mol dm−3 |
-sulfate-heptahydrate-sample.jpg)
As pH rises above 0 the above yellow hydrolyzed species form and as it rises above 2–3, reddish-brown hydrousiron(III) oxide precipitates out of solution. Although Fe3+ has a d5 configuration, its absorption spectrum is not like that of Mn2+ with its weak, spin-forbidden d–d bands, because Fe3+ has higher positive charge and is more polarizing, lowering the energy of its ligand-to-metalcharge transfer absorptions. Thus, all the above complexes are rather strongly colored, with the single exception of the hexaquo ion – and even that has a spectrum dominated by charge transfer in the near ultraviolet region.[16] On the other hand, the pale green iron(II) hexaquo ion[Fe(H2O)6]2+ does not undergo appreciable hydrolysis. Carbon dioxide is not evolved whencarbonate anions are added, which instead results in whiteiron(II) carbonate being precipitated out. In excess carbon dioxide this forms the slightly soluble bicarbonate, which occurs commonly in groundwater, but it oxidises quickly in air to formiron(III) oxide that accounts for the brown deposits present in a sizeable number of streams.[17]
Due to its electronic structure, iron has a very large coordination and organometallic chemistry.
Many coordination compounds of iron are known, although few exhibithigh-valent iron. A typical six-coordinate anion is hexachloroferrate(III), [FeCl6]3−, found in the mixedsalttetrakis(methylammonium) hexachloroferrate(III) chloride.[18][19] Complexes with multiple bidentate ligands havegeometric isomers. For example, thetrans-chlorohydridobis(bis-1,2-(diphenylphosphino)ethane)iron(II) complex is used as a starting material for compounds with theFe(dppe)2moiety.[20][21] The ferrioxalate ion with threeoxalate ligands (shown at right) displayshelical chirality with its two non-superposable geometries labelledΛ (lambda) for the left-handed screw axis andΔ (delta) for the right-handed screw axis, in line with IUPAC conventions.[16]Potassium ferrioxalate is used in chemicalactinometry and along with itssodium salt undergoesphotoreduction applied in old-style photographic processes. Thedihydrate ofiron(II) oxalate has apolymeric structure with co-planar oxalate ions bridging between iron centres with the water of crystallisation located forming the caps of each octahedron, as illustrated below.[22]
(H2O)2-chain-from-xtal-2008-CM-3D-balls.png)
iron(III)_chloride.jpg)
Iron(III) complexes are quite similar to those ofchromium(III) with the exception of iron(III)'s preference forO-donor instead ofN-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 forphenols orenols. For example, in theferric chloride test, used to determine the presence of phenols,iron(III) chloride reacts with a phenol to form a deep violet complex:[16]
Among the halide and pseudohalide complexes, fluoro complexes of iron(III) are the most stable, with the colorless [FeF5(H2O)]2− being the most stable in aqueous solution. Chloro complexes are less stable and favor tetrahedral coordination as in [FeCl4]−; [FeBr4]− and [FeI4]− are reduced easily to iron(II).Thiocyanate is a common test for the presence of iron(III) as it forms the blood-red [Fe(SCN)(H2O)5]2+. Like manganese(II), most iron(III) complexes are high-spin, the exceptions being those with ligands that are high in thespectrochemical series such ascyanide. An example of a low-spin iron(III) complex is [Fe(CN)6]3−. The cyanide ligands may easily be detached in [Fe(CN)6]3−, and hence this complex is poisonous, unlike the iron(II) complex [Fe(CN)6]4− found in Prussian blue,[16] which does not releasehydrogen cyanide except when dilute acids are added.[17] Iron shows a great variety of electronicspin states, including every possible spin quantum number value for a d-block element from 0 (diamagnetic) to5⁄2 (5 unpaired electrons). This value is always half the number of unpaired electrons. Complexes with zero to two unpaired electrons are considered low-spin and those with four or five are considered high-spin.[12]
Iron(II) complexes are less stable than iron(III) complexes but the preference forO-donor ligands is less marked, so that for example[Fe(NH3)6]2+ is known while[Fe(NH3)6]3+ is not. They have a tendency to be oxidized to iron(III) but this can be moderated by low pH and the specific ligands used.[17]
Organoiron chemistry is the study oforganometallic compounds of iron, where carbon atoms are covalently bound to the metal atom. They are many and varied, includingcyanide complexes,carbonyl complexes,sandwich andhalf-sandwich compounds.
Prussian blue or "ferric ferrocyanide", Fe4[Fe(CN)6]3, is an old and well-known iron-cyanide complex, extensively used as pigment and in several other applications. Its formation can be used as a simple wet chemistry test to distinguish between aqueous solutions of Fe2+ and Fe3+ as they react (respectively) withpotassium ferricyanide andpotassium ferrocyanide to form Prussian blue.[7]
Another old example of an organoiron compound isiron pentacarbonyl, Fe(CO)5, in which a neutral iron atom is bound to the carbon atoms of fivecarbon monoxide molecules. The compound can be used to makecarbonyl iron powder, a highly reactive form of metallic iron.Thermolysis of iron pentacarbonyl givestriiron dodecacarbonyl,Fe3(CO)12, a complex with a cluster of three iron atoms at its core. Collman's reagent,disodium tetracarbonylferrate, is a useful reagent for organic chemistry; it contains iron in the −2 oxidation state.Cyclopentadienyliron dicarbonyl dimer contains iron in the rare +1 oxidation state.[23]
.JPG)
A landmark in this field was the discovery in 1951 of the remarkably stablesandwich compoundferroceneFe(C5H5)2, by Pauson and Kealy[24] and independently by Miller and colleagues,[25] whose surprising molecular structure was determined only a year later byWoodward andWilkinson[26] andFischer.[27]Ferrocene is still one of the most important tools and models in this class.[28]
Iron-centered organometallic species are used ascatalysts. TheKnölker complex, for example, is atransfer hydrogenation catalyst forketones.[29]
The iron compounds produced on the largest scale in industry areiron(II) sulfate (FeSO4·7H2O) andiron(III) chloride (FeCl3). The former is one of the most readily available sources of iron(II), but is less stable to aerial oxidation thanMohr's salt ((NH4)2Fe(SO4)2·6H2O). Iron(II) compounds tend to be oxidized to iron(III) compounds in the air.[7]
