Iron(III) chloride describes the inorganic compounds with the formula FeCl3(H2O)x. Also called ferric chloride, these compounds are available both in anhydrous and hydrated forms which are both hygroscopic. They are common sources of iron in its +3 oxidation state. The anhydrous derivative is a Lewis acid, while the hydrate is a mild oxidizing agent. It is used as a water cleaner and as an etchant for metals.
Structure and properties
All forms of ferric chloride are paramagnetic, owing to the presence of five unpaired electrons residing in 3d orbitals. This electronic configuration places electrons in molecular orbitals that are antibonding with respect to ligands. Thus, iron(III) chlorides are labile, undergoing rapid ligand exchange in solution. In contrast to their kinetic lability, iron(III) chlorides are thermodynamically robust, as reflected by the vigorous methods applied to their synthesis.
The anhydrous compound is a hygroscopic crystalline solid with a melting point of 307.6 °C. The colour depends on the viewing angle: by reflected light, the crystals appear dark green, but by transmitted light, they appear purple-red. Anhydrous iron(III) chloride has the BiI3 structure, with octahedral Fe(III) centres interconnected by two-coordinate chloride ligands.
In addition to the anhydrous material, ferric chloride aggressively forms hydrates upon exposure to water, reflecting its Lewis acidity. Four of these hydrates have been crystallized and examined by X-ray crystallography. They all exhibit deliquesce and feature trans-[FeCl2(H2O)4]+ cations, with either chloride or [FeCl4]− as the anions.
dihydrate: FeCl3·2H2O has the structural formula trans-[FeCl2(H2O)4][FeCl4].
FeCl3·2.5H2O has the structural formula cis-[FeCl2(H2O)4][FeCl4]·H2O.
FeCl3·3.5H2O has the structural formula cis-[FeCl2(H2O)4][FeCl4]·3H2O.
hexahydrate: FeCl3·6H2O has the structural formula trans-[FeCl2(H2O)4]Cl·2H2O.
Aqueous solutions of ferric chloride are characteristically yellow, in contrast to the pale pink solutions of [Fe(H2O)6]3+. Thus, the chloride ligand significantly influences the optical properties of the iron center. According to spectroscopic measurements, the main species in aqueous solutions of ferric chloride are the octahedral [FeCl2(H2O)4]+ (stereochemistry unspecified) and the tetrahedral [FeCl4]−. The cationic aquo complex is strongly acidic:
[FeCl2(H2O)4)]+ ⇌ [FeCl2(OH)(H2O)3] + H+
Anhydrous iron(III) chloride dissolves in diethyl ether and tetrahydrofuran forming 1:2 adducts of the formula FeCl3(ether)2. In these complexes, the iron is pentacoordinate.
Several hundred thousand kilograms of anhydrous iron(III) chloride are produced annually. The principal method called direct chlorination, uses scrap iron as a precursor:
2 Fe + 3 Cl2 → 2 FeCl3
The reaction is conducted at several hundred degrees such that the product is gaseous. Using excess chlorine guarantees that the intermediate ferrous chloride is converted to the ferric state. A similar but laboratory-scale process also has been described.
Solutions of iron(III) chloride are produced industrially both from iron and from ore, in a closed-loop process.
Heating hydrated iron(III) chloride does not yield anhydrous ferric chloride. Instead, the solid decomposes into hydrochloric acid and iron oxychloride. Hydrated iron(III) chloride can be converted to the anhydrous form by treatment with thionyl chloride. Similarly, dehydration can be done with trimethylsilyl chloride:
Alkali metal alkoxides react to give the iron(III) alkoxide complexes. These products have more complicated structures than anhydrous iron(III) chloride. In the solid phase a variety of multinuclear complexes have been described for the nominal stoichiometric reaction between FeCl3 and sodium ethoxide:
In addition to these simple stoichiometric reactions, the Lewis acidity of ferric chloride enables its use in a variety of acid-catalyzed reactions as described below in the section on organic chemistry.
Iron(III) chloride is a oxidizing agent. A simple illustration is its tendency to release chlorine when heated above 160 °C:
To a significant extent, iron(III) acetylacetonate and related beta-diketonate complexes are more widely used than FeCl3 as ether-soluble sources of ferric ion. These diketonate complexes have the advantages that they do not form hydrates, unlike iron(III) chloride, and they are more soluble in relevant solvents.Cyclopentadienyl magnesium bromide undergoes a complex reaction with iron(III) chloride, resulting in ferrocene:
As a reagent in organic chemistry, iron(III) chloride has attracted interest for both its redox activity and its Lewis acidity. Furthermore, because they are inexpensive and relatively nontoxic, iron chlorides have been widely examined. Illustrating it use as a Lewis acid, iron(III) chloride catalyseselectrophilic aromatic substitution and chlorinations. In this role, its function is similar to that of aluminium chloride. In some cases, mixtures of the two are used. Iron(III) chloride oxidizes naphthols to naphthoquinones:
Anhydrous iron(III) chloride is harmful, highly corrosive, and acidic. Soluble iron(III) salts irritate the eyes, skin, mucous membrane, and induce abdominal pain, diarrhea, and vomiting when ingested.
FeCl3 naturally occurs as the rare mineral molysite, usually related to volcanic and other-type fumaroles.
FeCl3 is also produced as an atmospheric salt aerosol by a reaction between iron-rich dust and hydrochloric acid from sea salt. This iron salt aerosol causes about 5% of naturally-occurring oxidization of methane and is thought to have a range of cooling effects.
^An alternative GHS classification from the Japanese GHS Inter-ministerial Committee (2006) notes the possibility of respiratory tract irritation from FeCl3 and differs slightly in other respects from the classification used here.
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