Iron(III) chloride

Iron(III) chloride (hydrate)
Iron(III) chloride (anhydrous)
Names
IUPAC names Iron(III) chloride Iron trichloride
Other names Ferric chloride Molysite Flores martis
Identifiers
CAS Number	7705-08-0 Y 10025-77-1 (hexahydrate) N
3D model (JSmol)	Interactive image
ChEBI	CHEBI:30808 Y
ChemSpider	22792 Y
ECHA InfoCard	100.028.846
EC Number	231-729-4
PubChem CID	24380
RTECS number	LJ9100000
UNII	U38V3ZVV3V N 0I2XIN602U (hexahydrate) N
UN number	1773 (anhydrous) 2582 (aq. soln.)
CompTox Dashboard (EPA)	DTXSID8020622
InChI InChI=1S/3ClH.Fe/h3*1H;/q;;;+3/p-3 Y Key: RBTARNINKXHZNM-UHFFFAOYSA-K Y InChI=1S/3ClH.Fe/h3*1H;/q;;;+3/p-3 Key: RBTARNINKXHZNM-DFZHHIFOAF Key: RBTARNINKXHZNM-UHFFFAOYSA-K
SMILES Cl[Fe](Cl)Cl
Properties
Chemical formula
Molar mass	162.204g/mol (anhydrous) 270.295g/mol (hexahydrate)[1]
Appearance	Green-black by reflected light; purple-red by transmitted light; yellow solid as hexahydrate; brown as aq. solution
Odor	Slight HCl
Density	2.90g/cm3 (anhydrous) 1.82g/cm3 (hexahydrate)[1]
Melting point	307.6 °C (585.7 °F; 580.8 K) (anhydrous) 37 °C (99 °F; 310 K) (hexahydrate)[1]
Boiling point	316 °C (601 °F; 589 K) (anhydrous, decomposes)[1] 280 °C (536 °F; 553 K) (hexahydrate, decomposes)
Solubility in water	912g/L (anh. or hexahydrate, 25°C)[1]
Solubility in Acetone Methanol Ethanol Diethyl ether[1]	630 g/L (18 °C) Highly soluble 830 g/L Highly soluble
Magnetic susceptibility (χ)	+13,450·10-6cm3/mol[2]
Viscosity	12 cP (40% solution)
Structure
Crystal structure	Hexagonal, hR24
Space group	R3, No. 148[3]
Lattice constant	a = 0.6065nm, b = 0.6065nm, c = 1.742nm ? = 90°, ? = 90°, ? = 120°
Formula units (Z)	6
Coordination geometry	Octahedral
Hazards[5][6][Note 1]
Safety data sheet	ICSC 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, P390, P405, P406, P501
NFPA 704 (fire diamond)	0 2 0
Flash point	Non-flammable
NIOSH (US health exposure limits):
REL (Recommended)	TWA 1mg/m3[4]
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
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YN ?)
Infobox references

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 D_3hpoint group molecular 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

• 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:^[10]

${\displaystyle {\ce {2{Fe_{(}s)}+3Cl2_{(}g)->2FeCl3_{(}s)}}}$

Solutions of iron(III) chloride are produced industrially both from iron and from ore, in a closed-loop process.

1. Dissolving iron ore in hydrochloric acid

   ${\displaystyle {\ce {Fe3O4_{(}s){+~}8HCl_{(}aq)->FeCl2_{(}aq){+~}2FeCl3_{(}aq){+~}4H2O_{(}l)}}}$

2. Oxidation of iron(II) chloride with chlorine

   ${\displaystyle {\ce {2FeCl2_{(}aq){+~}Cl2_{(}g)->2FeCl3_{(}aq)}}}$

3. Oxidation of iron(II) chloride with oxygen

   ${\displaystyle {\ce {4FeCl2_{(}aq){+~}O2{+~}4HCl->4FeCl3_{(}aq){+~}2H2O_{(}l)}}}$

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.^[11] Similarly, dehydration can be effected with trimethylsilyl chloride:^[12]

${\displaystyle {\ce {FeCl3.6H2O + 12 Me3SiCl -> FeCl3 + 6 (Me3Si)2O + 12 HCl}}}$

Reactions

A brown, acidic solution of iron(III) chloride

Iron(III) chloride undergoes hydrolysis to give a strongly acidic solution.^[13][9]

When heated with iron(III) oxide at 350 °C, iron(III) chloride gives iron oxychloride, a layered solid and intercalation host.^[14]

${\displaystyle {\ce {FeCl3 + Fe2O3 -> 3FeOCl}}}$

The anhydrous salt is a moderately strong Lewis acid, forming adducts with Lewis bases such as triphenylphosphine oxide; e.g., where Ph is phenyl. It also reacts with other chloride salts to give the yellow tetrahedral ion. Salts of in hydrochloric acid can be extracted into diethyl ether.

Redox reactions

Iron(III) chloride is a mild oxidising agent, for example, it is capable of oxidising copper(I) chloride to copper(II) chloride.

${\displaystyle {\ce {FeCl3 + CuCl -> FeCl2 + CuCl2}}}$

It also reacts with iron to form iron(II) chloride:

${\displaystyle {\ce {2 FeCl3 + Fe -> 3 FeCl2}}}$

A traditional synthesis of anhydrous ferrous chloride is the reduction of FeCl₃ with chlorobenzene:^[15]

${\displaystyle {\ce {2 FeCl3 + C6H5Cl -> 2 FeCl2 + C6H4Cl2 + HCl}}}$

With carboxylate anions

Oxalates react rapidly with aqueous iron(III) chloride to give . Other carboxylate salts form complexes; e.g., citrate and tartrate.

With alkali metal alkoxides

Alkali metal alkoxides react to give the metal alkoxide complexes of varying complexity.^[16] The compounds can be dimeric or trimeric.^[17] In the solid phase a variety of multinuclear complexes have been described for the nominal stoichiometric reaction between and sodium ethoxide:^[18][19]

${\displaystyle {\ce {FeCl3 + 3 [C2H5O]- Na+ -> Fe(OC2H5)3 + 3 NaCl}}}$

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) :^[20][21]

${\displaystyle {\ce {2 FeCl3 + LiCH3 -> FeCl2 + LiFeCl4 + .CH3}}}$

${\displaystyle {\ce {LiFeCl4 + LiCH3 -> Li2FeCl4 + .CH3}}}$

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.^[22] 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.

${\displaystyle {\ce {{[Fe(H2O)6]^{3+}}+4HO^{-}->{[Fe(H2O)2(HO)4]^{-}}+4H2O->{[Fe(H2O)O(HO)2]^{-}}+6H2O}}}$

It is also used as a leaching agent in chloride hydrometallurgy,^[23] for example in the production of Si from FeSi (Silgrain process).^[24]

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.^[25]

${\displaystyle {\ce {FeCl3 + Cu -> FeCl2 + CuCl}}}$

${\displaystyle {\ce {FeCl3 + CuCl -> FeCl2 + CuCl2}}}$

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.

${\displaystyle {\ce {H2C=CH2 + Cl2 -> ClCH2CH2Cl}}}$

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.^[26] 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.

This test can be used to detect the presence of gamma-hydroxybutyric acid and gamma-butyrolactone,^[27] which cause it to turn red-brown.

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 in wastewater treatment for odor control.
• 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) by etching copper.
• Used to strip aluminium coating from mirrors.
• Used to etch intricate medical devices.
• 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.^[28]
• 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.^[29]
• Used in an animal thrombosis model.^[30]
• Used in energy storage systems.
• Historically it was used to make direct positive blueprints.^[31][32]
• A component of modified Carnoy's solution used for surgical treatment of keratocystic odontogenic tumor (KOT).

Safety

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.

See also

• Verhoeff's stain

Notes

References

Further reading

1. Lide DR, ed. (1990). CRC Handbook of Chemistry and Physics (71st ed.). Ann Arbor, MI, USA: CRC Press. ISBN 9780849304712.
2. Stecher PG, Finkel MJ, Siegmund OH, eds. (1960). The Merck Index of Chemicals and Drugs (7th ed.). Rahway, NJ, USA: Merck & Co.
3. Nicholls D (1974). Complexes and First-Row Transition Elements, Macmillan Press, London, 1973. A Macmillan chemistry text. London: Macmillan Press. ISBN 9780333170885.
4. Wells AF (1984). Structural Inorganic Chemistry. Oxford science publications (5th ed.). Oxford, UK: Oxford University Press. ISBN 9780198553700.
5. March J (1992). Advanced Organic Chemistry (4th ed.). New York: John Wiley & Sons, Inc. pp. 723. ISBN 9780471581482.
6. Reich HJ, Rigby HJ, eds. (1999). Acidic and Basic Reagents. Handbook of Reagents for Organic Synthesis. New York: John Wiley & Sons, Inc. ISBN 9780471979258.