Ferric Chloride

Iron(III) chloride
Identifiers
CAS number 7705-08-0 YesY
10025-77-1 (hexahydrate)
PubChem 24380
ChemSpider 22792 YesY
UN number 1773 (anhydrous)
2582 (aq. soln.)
ChEBI CHEBI:30808 YesY
RTECS number LJ9100000
Jmol-3D images Image 1
Properties
Molecular formula FeCl3
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
Odor slight HCl
Density 2.898 g/cm3 (anhydrous)
1.82 g/cm3 (hexahydrate)
Melting point

306 °C (anhydrous)
37 °C (hexahydrate)

Boiling point

315 °C (anhydrous, decomp)
280 °C (hexahydrate, decomp) (partial decomposition to FeCl2 + Cl2)

Solubility in water 74.4 g/100 mL (0 °C)[1]
92 g/100 mL (hexahydrate, 20 °C)
Solubility in acetone
Methanol
Ethanol
Diethyl ether
63 g/100 ml (18 °C)
highly soluble
83 g/100 ml
highly soluble
Viscosity 40% solution: 12 cP
Structure
Crystal structure hexagonal
Coordination
geometry
octahedral
Hazards[2][3][Note 1]
MSDS 1499
GHS pictograms
GHS signal word DANGER
GHS hazard statements H290, H302, H314, H318
GHS precautionary statements P234, P260, P264, P270, P273, P280, P301+312, P301+330+331, P303+361+353, P363, P304+340, P310, P321, P305+351+338
EU Index listed
NFPA 704
0
2
0
Flash point non-flammable
Related compounds
Other anions Iron(III) fluoride
Iron(III) bromide
Other cations Iron(II) chloride
Manganese(II) chloride
Cobalt(II) chloride
Ruthenium(III) chloride
Related coagulants Iron(II) sulfate
Polyaluminium chloride
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Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

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.[5]

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.[6]

Preparation

Anhydrous iron(III) chloride may be prepared by union of the elements:[7]

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.

  1. Dissolving pure iron in a solution of iron(III) chloride
    Fe(s) + 2 FeCl3(aq) → 3 FeCl2(aq)
  2. Dissolving iron ore in hydrochloric acid
    Fe3O4(s) + 8 HCl(aq) → FeCl2(aq) + 2 FeCl3(aq) + 4 H2O
  3. Oxidation of iron (II) chloride with chlorine
    2 FeCl2(aq) + Cl2(g) → 2 FeCl3(aq)
  4. Oxidation of iron (II) chloride with oxygen
    FeCl2(aq) + ¼O2 + HCl → FeCl3(aq) + ½H2O

Like many other hydrated metal chlorides, hydrated iron(III) chloride can be converted to the anhydrous salt by refluxing with thionyl chloride.[8] 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.[9] The compounds can be dimeric or trimeric.[10] In the solid phase a variety of multinuclear complexes have been described for the nominal stoichiometric reaction between FeCl3 and sodium ethoxide:[11][12]

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.[13] 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,[14] for example in the production of Si from FeSi. (Silgrain process)[15]

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.[16]

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

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.[17] 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.

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 (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.[18]
  • 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.[19]
  • Used in an animal thrombosis model.[20]

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

Notes and references

Notes

References

Further reading

  1. Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
  2. The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
  3. D. Nicholls, Complexes and First-Row Transition Elements, Macmillan Press, London, 1973.
  4. A.F. Wells, 'Structural Inorganic Chemistry, 5th ed., Oxford University Press, Oxford, UK, 1984.
  5. J. March, Advanced Organic Chemistry, 4th ed., p. 723, Wiley, New York, 1992.
  6. Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents, (H. J. Reich, J. H. Rigby, eds.), Wiley, New York, 1999.
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