Iron(III) chloride, also 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. 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.
Nomenclature 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 which increasingly dissociates into the monomeric FeCl3 at higher
temperature, in competition with its reversible decomposition to give iron(II) chloride and
chlorine gas. Preparation
Anhydrous iron(III) chloride may be prepared by union of the elements:
2 Fe(s) + 3 Cl2(g) → 2 FeCl3(s) 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
Oxidation of iron chloride with chlorine 2 FeCl2(aq) + Cl2(g) → 2 FeCl3(aq)
Oxidation of iron chloride with oxygen FeCl2(aq) + ¼O2 + HCl → FeCl3(aq) + ½H2O
Reacting Iron with hydrochloric acid, then with hydrogen peroxide. The hydrogen peroxide
is the catalyst in turning iron 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. Reactions
Iron(III) chloride undergoes hydrolysis to give an acidic solution. When heated with
iron(III) oxide at 350 °C, iron(III) chloride gives iron oxychloride, a layered solid and
intercalation host. 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
Oxalates react rapidly with aqueous iron(III) chloride to give [Fe(C2O4)3]3−. Other carboxylate
salts form complexes, e.g. citrate and tartrate. Oxidation
Iron(III) chloride is a mild oxidising agent, for example, it is capable of oxidising copper(I)
chloride to copper(II) chloride. 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).
Uses Industrial
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 It is also used as a leaching agent in chloride
hydrometallurgy, for example in the production of Si from FeSi.
Another important application of iron(III) chloride is etching copper in two-step redox
reaction to copper(I) chloride and then to copper(II) chloride in the production of printed
circuit boards. 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, an important commodity
chemical, which is mainly used for the industrial production of vinyl chloride, the monomer
for making PVC. H2C=CH2 + Cl2 → ClCH2CH2Cl
Laboratory use 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 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. Other uses
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 knife craftsmen and sword smiths to stain blades, as to give a contrasting effect
to the metal, and 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. 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.
Safety 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.
See also Iron(II) sulfate
Aluminium chloride Verhoeff’s Stain
Notes and references Notes
References
Further reading 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,, Wiley, New York,
1999.