|IUPAC name||Iron(III) chloride|
|Other names||ferric chloride|
molysite mineral form
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|Molar mass||162.20 g·mol−1 |
hexahydrate: 270.3 g·mol−1
|Appearance||green-black by reflected light; purple-red by transmitted light|
hexahydrate: yellow solid
aq. solutions: brown
40% solution: 1.4 g·ml−1
|Solubility in acetone
|63 g/100 ml (18 °C)|
83 g/100 ml
|Viscosity||40% solution: 12 cP|
|Main hazards||Very corrosive|
|Except where noted otherwise, data are given for|
materials in their standard state
(at 25 °C, 100 kPa)
Infobox disclaimer and references
Iron(III) chloride, generically called ferric chloride, is an industrial scale commodity chemical compound, with the formula FeCl3. The colour of iron(III) chloride crystals depends on the viewing angle: by reflected light the crystals appear dark green, but by transmitted light they appear purple-red. Iron(III) chloride is deliquescent, fuming in moist air due to the evolution of HCl, which hydrates, giving a mist.
When dissolved in water, iron(III) chloride undergoes hydrolysis and gives off heat in an exothermic reaction. The resulting brown, acidic, and corrosive solution is used as a coagulant in sewage treatment and drinking water production, and as an etchant for copper-based metals in printed circuit boards. Anhydrous iron(III) chloride is a fairly strong Lewis acid, and it is used as a catalyst in organic synthesis.
Chemical and physical properties
Iron(III) chloride has a relatively low melting point and boils at around 315 °C. The vapour consists of the dimer Fe2Cl6 (compare aluminium chloride) which increasingly dissociates into the monomeric FeCl3 (D3h point group molecular symmetry) at higher temperature, in competition with its reversible decomposition to give iron(II) chloride and chlorine gas.
- FeCl3 + Fe2O3 → 3 FeOCl
Iron(III) chloride is a mild oxidising agent, for example capable of oxidising copper(I) chloride to copper(II) chloride. Reducing agents such as hydrazine convert iron(III) chloride to complexes of iron(II).
Iron(III) chloride adopts the BI3 structure, which features octahedral Fe(III) centres interconnected by two-coordinate chloride ligands.
Preparation and production
Anhydrous iron(III) chloride may be prepared by union of the elements:
Solutions of iron(III) chloride are produced industrially both from iron and from ore, in a closed-loop process.
- Dissolving pure iron in a solution of iron(III) chloride
Fe(s) + 2 FeCl3(aq) → 3 FeCl2(aq)
- Dissolving iron ore in hydrochloric acid
Fe3O4(s) + 8 HCl(aq) → FeCl2(aq) + 2 FeCl3(aq) + 4 H2O
- Upgrading the iron(II) chloride with chlorine
2 FeCl2(aq) + Cl2(g) → 2 FeCl3(aq)
Alternatively, iron(II) chloride can be oxidised with sulfur dioxide:
- 32 FeCl2 + 8 SO2 + 32 HCl → 32 FeCl3 + S8 + 16 H2O
Hydrated iron(III) chloride can be converted to the anhydrous salt by heating with thionyl chloride. The hydrate cannot be converted to anhydrous iron(III) chloride by only heat, as instead HCl is evolved and iron oxychloride forms.
In industrial application, iron(III) chloride is used in sewage treatment and drinking water production, where FeCl3 in slightly basic water reacts with the hydroxide ion to form a floc of iron(III) hydroxide, or more correctly formulated as FeO(OH)-, that can remove suspended materials.
- Fe3+ + 4 OH- → Fe(OH)4- → FeO(OH)2-.H2O
- FeCl3 + Cu → FeCl2 + CuCl
- FeCl3 + CuCl → FeCl2 + CuCl2
Iron(III) chloride is used as catalyst for the reaction of ethylene with chlorine, forming ethylene dichloride (1,2-dichloroethane), an important commodity chemical, which is mainly used for the industrial production of vinyl chloride, the monomer for making PVC.
- H2C=CH2 + Cl2 → ClCH2CH2Cl
In the laboratory iron(III) chloride is commonly employed as a Lewis acid for catalysing reactions such as chlorination of aromatic compounds and Friedel-Crafts reaction of aromatics. It is less powerful than aluminium chloride, but in some cases this mildness leads to higher yields, for example in the alkylation of benzene:
The "ferric chloride test" is a traditional colorimetric test for phenols, which uses a 1% iron(III) chloride solution that has been neutralised with sodium hydroxide until a slight precipitate of FeO(OH) is formed. The mixture is filtered before use. The organic substance is dissolved in water, methanol or ethanol, then the neutralised iron(III) chloride solution is added—a transient or permanent coloration (usually purple, green or blue) indicates the presence of a phenol or enol.
Iron(III) chloride is sometimes used by American coin collectors to identify the dates of Buffalo nickels that are so badly worn that the date is no longer visible.
Iron(III) chloride is commonly used by knife craftsmen and sword smiths to stain blades, as to give a contrasting effect to the metal, and to view metal layering or imperfections.
Iron(III) chloride is also used in veterinary practice to treat overcropping of an animal's claws, particularly when the overcropping results in bleeding.
Iron(III) chloride is toxic, highly corrosive and acidic. The anhydrous material is a powerful dehydrating agent. In secondary/high schools all around the world, where Design and technology is a subject taught, Ferric Chloride used for PCB etching is usually diluted with water. Despite this, hands and other surfaces that have contacted it should still be washed immediately after one finishes with it.
- Holleman, A.F. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5. Unknown parameter
- Tarr, B.R. (1950). "Anhydrous Iron(III) Chloride". Inorganic Syntheses. III: pp 191–194.
- Greenwood, N.N. (1997). Chemistry of the Elements (2nd ed. ed.). Oxford: Butterworth-Heinemann. Unknown parameter
- Furnell, B.S. (1989). Vogel's Textbook of Practical Organic Chemistry (5th edition ed.). New York: Longman/Wiley. Unknown parameter
- Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
- The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
- D. Nicholls, Complexes and First-Row Transition Elements, Macmillan Press, London, 1973.
- A.F. Wells, 'Structural Inorganic Chemistry, 5th ed., Oxford University Press, Oxford, UK, 1984.
- J. March, Advanced Organic Chemistry, 4th ed., p. 723, Wiley, New York, 1992.
- Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents, (H. J. Reich, J. H. Rigby, eds.), Wiley, New York, 1999.