|Jmol 3D model||Interactive image|
|UN number||1773 (anhydrous)
2582 (aq. soln.)
|Molar mass||162.2 g/mol (anhydrous)
270.3 g/mol (hexahydrate)
|Appearance||green-black by reflected light; purple-red by transmitted light
hexahydrate: yellow solid
aq. solutions: brown
|Density||2.898 g/cm3 (anhydrous)
1.82 g/cm3 (hexahydrate)
|Melting point||306 °C (583 °F; 579 K) (anhydrous)
37 °C (99 °F; 310 K) (hexahydrate)
|Boiling point||315 °C (599 °F; 588 K) (anhydrous, decomposes)
280 °C (536 °F; 553 K) (hexahydrate, decomposes) partial decomposition to FeCl2 + Cl2
|74.4 g/100 mL (0 °C)
92 g/100 mL (hexahydrate, 20 °C)
|Solubility in acetone
|63 g/100 ml (18 °C)
83 g/100 ml
|Viscosity||40% solution: 12 cP|
|Safety data sheet||ICSC|
|GHS signal word||DANGER|
|H290, H302, H314, H318|
|P234, P260, P264, P270, P273, P280, P301+312, P301+330+331, P303+361+353, P363, P304+340, P310, P321, P305+351+338|
|US health exposure limits (NIOSH):|
|TWA 1 mg/m3|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Iron(III) chloride, also called ferric chloride, is an industrial scale commodity chemical compound, with the formula FeCl3 and with iron in the +3 oxidation state. 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. Anhydrous iron(III) chloride is deliquescent, forming hydrated hydrogen chloride mists in moist air. It is rarely observed in its natural form, mineral molysite, known mainly from some fumaroles.
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 flocculant 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.
The descriptor hydrated or anhydrous is used when referring to iron(III) chloride, to distinguish between the two common forms. The hexahydrate is usually given as the simplified empirical formula FeCl3⋅6H2O. It may also be given as trans-[Fe(H2O)4Cl2]Cl⋅2H2O and the systematic name tetraaquadichloroiron(III) chloride dihydrate, which more clearly represents its structure.
Structure and properties
Anhydrous iron(III) chloride adopts the BiI3 structure, which features octahedral Fe(III) centres interconnected by two-coordinate chloride ligands. Iron(III) chloride hexahydrate consists of trans-[Fe(H2O)4Cl2]+ cationic complexes and chloride anions, with the remaining two H2O molecules embedded within the monoclinic crystal structure.
Iron(III) chloride has a relatively low melting point and boils at around 315 °C. The vapour consists of the dimer Fe2Cl6 (c.f. 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.
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 iron ore in hydrochloric acid
- Fe3O4(s) + 8 HCl(aq) → FeCl2(aq) + 2 FeCl3(aq) + 4 H2O
- Oxidation of iron (II) chloride with chlorine
- 2 FeCl2(aq) + Cl2(g) → 2 FeCl3(aq)
- Oxidation of iron (II) chloride with oxygen
- 4FeCl2(aq) + O2 + 4HCl → 4FeCl3(aq) + 2H2O
- Reacting Iron with hydrochloric acid, then with hydrogen peroxide. The hydrogen peroxide is the oxidant in turning ferrous chloride into ferric chloride
Like many other hydrated metal chlorides, hydrated iron(III) chloride can be converted to the anhydrous salt by refluxing with thionyl chloride. Conversion of the hydrate to anhydrous iron(III) chloride is not accomplished by heating, as HCl and iron oxychlorides are produced.
- FeCl3 + Fe2O3 → 3 FeOCl
It is a moderately strong Lewis acid, forming adducts with Lewis bases such as triphenylphosphine oxide, e.g. FeCl3(OPPh3)2 where Ph = phenyl. It also reacts with other chloride salts to give the yellow tetrahedral FeCl4− ion. Salts of FeCl4− in hydrochloric acid can be extracted into diethyl ether.
Alkali metal alkoxides react to give the metal alkoxide complexes of varying complexity. The compounds can be dimeric or trimeric. In the solid phase a variety of multinuclear complexes have been described for the nominal stoichiometric reaction between FeCl3 and sodium ethoxide:
- FeCl3 + 3 [C2H5O]−Na+ → Fe(OC2H5)3 + 3 NaCl
- FeCl3 + CuCl → FeCl2 + CuCl2
It also reacts with iron to form iron(II) chloride:
- 2 FeCl3 + Fe → 3 FeCl2
Reducing agents such as hydrazine convert iron(III) chloride to complexes of iron(II).
In industrial application, iron(III) chloride is used in sewage treatment and drinking water production. In this application, FeCl3 in slightly basic water reacts with the hydroxide ion to form a floc of iron(III) hydroxide, or more precisely formulated as FeO(OH)−, that can remove suspended materials.
- [Fe(H2O)6]3+ + 4 HO− → [Fe(H2O)2(HO)4]− + 4 H2O → [Fe(H2O)O(HO)2]− + 6 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.
- Used in anhydrous form as a drying reagent in certain reactions.
- Used to detect the presence of phenol compounds in organic synthesis e.g.: examining purity of synthesised Aspirin.
- Used in water and wastewater treatment to precipitate phosphate as iron(III) phosphate.
- Used by American coin collectors to identify the dates of Buffalo nickels that are so badly worn that the date is no longer visible.
- Used by bladesmiths and artisans in pattern welding to etch the metal, giving it a contrasting effect, to view metal layering or imperfections.
- Used to etch the widmanstatten pattern in iron meteorites.
- 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.
- Used to make printed circuit boards (PCBs).
- Used in veterinary practice to treat overcropping of an animal's claws, particularly when the overcropping results in bleeding.
- Reacts with cyclopentadienylmagnesium bromide in one preparation of ferrocene, a metal-sandwich complex.
- Sometimes used in a technique of Raku ware firing, the iron coloring a pottery piece shades of pink, brown, and orange.
- Used to test the pitting and crevice corrosion resistance of stainless steels and other alloys.
- Used in conjunction with NaI in acetonitrile to mildly reduce organic azides to primary amines.
- Used in an animal thrombosis model.
- Used in energy storage systems
- Historically it was used to make direct positive blueprints; U.S. patent 241,713, May 17, 1881
- A component of modified Carnoy's solution used for surgical treatment of keratocystic odontogenic tumor (KOT)
Iron(III) chloride is toxic, 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.
Notes and references
- 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.
- Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
- "NIOSH Pocket Guide to Chemical Hazards #0346". National Institute for Occupational Safety and Health (NIOSH).
- HSNO Chemical Classification Information Database, New Zealand Environmental Risk Management Authority, retrieved 2010-09-19
- Various suppliers, collated by the Baylor College of Dentistry, Texas A&M University. (accessed 2010-09-19)
- GHS classification – ID 831, Japanese GHS Inter-ministerial Committee, 2006, retrieved 2010-09-19
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- The synthesis of iron (III) ethoxide revisited: Characterization of the metathesis products of iron (III) halides and sodium ethoxide, Gulaim A. Seisenbaevaa, Suresh Gohila, Evgeniya V. Suslovab, Tatiana V. Rogovab, Nataliya Ya. Turovab, Vadim G. Kesslera, Inorganica Chimica Acta, Volume 358, Issue 12, 1/8/2005, pp. 3506–3512, online link
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- Tseng, Michael; Dozier, A.; Haribabu, B.; Graham, U. M. (2006). "Transendothelial migration of ferric ion in FeCl3 injured murine common carotid artery". Thrombosis Research. 118 (2): 275–280. doi:10.1016/j.thromres.2005.09.004. PMID 16243382.
- Lietze, Ernst (1888). Modern Heliographic Processes. New York: D. Van Norstrand Company. p. 65.
- 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.
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- J. March, Advanced Organic Chemistry, 4th ed., p. 723, Wiley, New York, 1992.
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