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Iron(III) chloride is the inorganic compound with the formula . Also called ferric chloride, it is a common compound of iron in the +3 oxidation state. The anhydrous compound is a crystalline solid with a melting point of 307.6 °C. The color depends on the viewing angle: by reflected light the crystals appear dark green, but by transmitted light they appear purple-red.
Structure and properties
Anhydrous
Anhydrous iron(III) chloride has the BiI3 structure, with octahedral Fe(III) centres interconnected by two-coordinate chloride ligands.[3]
Iron(III) chloride has a relatively low melting point and boils at around 315 °C. The vapour consists of the dimer (cf. aluminium chloride) which increasingly dissociates into the monomeric (with D3hpoint groupmolecular symmetry) at higher temperature, in competition with its reversible decomposition to give iron(II) chloride and chlorine gas.[8]
Hydrates
In addition to the anhydrous material, ferric chloride forms four hydrates. All forms of iron(III) chloride feature two or more chlorides as ligands, and three hydrates feature FeCl4-.[9]
hexahydrate: FeCl3.6H2O has the structural formula trans-[Fe(H2O)4Cl2]Cl.2H2O[10]
FeCl3.2.5H2O has the structural formula cis-[Fe(H2O)4Cl2][FeCl4].H2O.
dihydrate: FeCl3.2H2O has the structural formula trans-[Fe(H2O)4Cl2][FeCl4].
FeCl3.3.5H2O has the structural formula cis-[FeCl2(H2O)4][FeCl4].3H2O.
Aqueous solution
Aqueous solutions of ferric chloride are characteristically yellow, in contrast to the pale pink solutions of [Fe(H2O)6]3+. According to spectroscopic measurements, the main species in aqueous solutions of ferric chloride are the octahedral complex [FeCl2(H2O)4]+ (stereochemistry unspecified) and the tetrahedral [FeCl4]-.[9]
Preparation
Anhydrous iron(III) chloride may be prepared by union of the elements:[11]
Solutions of iron(III) chloride are produced industrially both from iron and from ore, in a closed-loop process.
Small amounts can be produced by reacting iron with hydrochloric acid, then with hydrogen peroxide. The hydrogen peroxide is the oxidant in turning ferrous chloride into ferric chloride
Anhydrous iron(III) chloride cannot be obtained from the hydrate by heating. Instead, the solid decomposes into HCl and iron oxychloride. The conversion can be accomplished by treatment with thionyl chloride.[12] Similarly, dehydration can be effected with trimethylsilyl chloride:[13]
Reactions
A brown, acidic solution of iron(III) chloride
Iron(III) chloride undergoes hydrolysis to give a strongly acidic solution.[14][9]
Alkali metal alkoxides react to give the metal alkoxide complexes of varying complexity.[17] The compounds can be dimeric or trimeric.[18] In the solid phase a variety of multinuclear complexes have been described for the nominal stoichiometric reaction between and sodium ethoxide:[19][20]
With organometallic compounds
Iron(III) chloride in ether solution oxidizes methyl lithium to give first light greenish yellow lithium tetrachloroferrate(III) solution and then, with further addition of methyl lithium, lithium tetrachloroferrate(II) :[21][22]
The methyl radicals combine with themselves or react with other components to give mostly ethane and some methane .
Uses
Industrial
Iron(III) chloride is used in sewage treatment and drinking water production as a coagulant and flocculant.[23] In this application, in slightly basic water reacts with the hydroxide ion to form a floc of iron(III) hydroxide, or more precisely formulated as , that can remove suspended materials.
It is also used as a leaching agent in chloride hydrometallurgy,[24] for example in the production of Si from FeSi (Silgrain process).[25]
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.
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.[27] 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.
This reaction is exploited in the Trinder spot test, which is used to indicate the presence of salicylates, particularly salicylic acid, which contains a phenolic OH group.
Necessary for the etching of photogravure plates for printing photographic and fine art images in intaglio and for etching rotogravure cylinders used in the printing industry.
Iron(III) chloride is harmful, highly corrosive and acidic. The anhydrous material is a powerful dehydrating agent.
Although reports of poisoning in humans are rare, ingestion of ferric chloride can result in serious morbidity and mortality. Inappropriate labeling and storage lead to accidental swallowing or misdiagnosis. Early diagnosis is important, especially in seriously poisoned patients.
Natural occurrence
The natural counterpart of FeCl3 is the rare mineral molysite, usually related to volcanic and other-type fumaroles.[35][36]
^An alternative GHS classification from the Japanese GHS Inter-ministerial Committee (2006)[7] notes the possibility of respiratory tract irritation from and differs slightly in other respects from the classification used here.
^Tarr BR, Booth HS, Dolance A (1950). "Anhydrous Iron(III) Chloride". In Audrieth LF (ed.). Inorganic Syntheses. 3. McGraw-Hill Book Company, Inc. pp. 191-194. doi:10.1002/9780470132340.ch51. ISBN9780470132340.
^
Seisenbaeva GA, Gohil S, Suslova EV, et al. (2005). "The synthesis of iron (III) ethoxide revisited: Characterization of the metathesis products of iron (III) halides and sodium ethoxide". Inorg. Chim. Acta. 358 (12): 3506-3512. doi:10.1016/j.ica.2005.03.048.
^Berthold HJ, Spiegl HJ (1972). "Über die Bildung von Lithiumtetrachloroferrat(II) bei der Umsetzung von Eisen(III)-chlorid mil Lithiummethyl (1:1) in ätherischer Lösung". Z. Anorg. Allg. Chem. (in German). 391 (3): 193-202. doi:10.1002/zaac.19723910302.
^Berthold HJ, Spiegl HJ (1972). "Über die Bildung von Lithiumtetrachloroferrat(II) bei der Umsetzung von Eisen(III)-chlorid mil Lithiummethyl (1:1) in ätherischer Lösung". Z. Anorg. Allg. Chem. (in German). 391 (3): 193-202. doi:10.1002/zaac.19723910302.
^Park KH, Mohapatra D, Reddy BR (2006). "A study on the acidified ferric chloride leaching of a complex (Cu-Ni-Co-Fe) matte". Separation and Purification Technology. 51 (3): 332-337. doi:10.1016/j.seppur.2006.02.013.
^Dueñas Díez M, Fjeld M, Andersen E, et al. (2006). "Validation of a compartmental population balance model of an industrial leaching process: The Silgrain process". Chem. Eng. Sci.61 (1): 229-245. doi:10.1016/j.ces.2005.01.047.
^
Zhang SY, Huang ZP (2006). "A color test for rapid screening of gamma-hydroxybutyric acid (GHB) and gamma-butyrolactone (GBL) in drink and urine". Fa Yi Xue Za Zhi. 22 (6): 424-7. PMID17285863.
^Kamal A, Ramana K, Ankati H, et al. (2002). "Mild and efficient reduction of azides to amines: synthesis of fused [2,1-b]quinazolines". Tetrahedron Lett.43 (38): 6861-6863. doi:10.1016/S0040-4039(02)01454-5.
Reich HJ, Rigby HJ, eds. (1999). Acidic and Basic Reagents. Handbook of Reagents for Organic Synthesis. New York: John Wiley & Sons, Inc. ISBN9780471979258.