Chemistry:Iron(III) chloride

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chemical compound
Iron(III) chloride
Iron(III) chloride hexahydrate.jpg
Iron(III) chloride (hydrate)
Iron(III) chloride anhydrate.jpg
Iron(III) chloride (anhydrous)
IUPAC names
Iron(III) chloride
Iron trichloride
Other names
  • Ferric chloride
  • Molysite
  • Flores martis
3D model (JSmol)
EC Number
  • 231-729-4
RTECS number
  • LJ9100000
UN number
  • 1773 (anhydrous)
  • 2582 (aq. soln.)
Molar mass
  • 162.204 g/mol (anhydrous)
  • 270.295 g/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
  • 2.90 g/cm3 (anhydrous)
  • 1.82 g/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)
912 g/L (anh. or hexahydrate, 25 °C)[1]
Solubility in
  • 630 g/L (18 °C)
  • Highly soluble
  • 830 g/L
  • Highly soluble
+13,450·10−6 cm3/mol[2]
Viscosity 12 cP (40% solution)
Hexagonal, hR24
R3, No. 148[3]
a = 0.6065 nm, b = 0.6065 nm, c = 1.742 nm
α = 90°, β = 90°, γ = 120°
Hazards[5][6][Note 1]
GHS pictograms Corr. Met. 1; Skin Corr. 1C; Eye Dam. 1Acute Tox. 4 (oral)
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, P390, P405, P406, P501
NFPA 704 (fire diamond)
Flammability code 0: Will not burn. E.g. waterHealth code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
NIOSH (US health exposure limits):
REL (Recommended)
TWA 1 mg/m3[4]
Related compounds
Other anions
Other cations
Related coagulants
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references
Tracking categories (test):

Iron(III) chloride is the inorganic compound with the formula (FeCl
). 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 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 Fe
(cf. aluminium chloride) which increasingly dissociates into the monomeric FeCl
(with D3h point group molecular symmetry) at higher temperature, in competition with its reversible decomposition to give iron(II) chloride and chlorine gas.[8]


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]


Anhydrous iron(III) chloride may be prepared by treating iron with chlorine:[11]

[math]\ce{ 2 {Fe_(s)} + 3 Cl2_(g) -> 2 FeCl3_(s) }[/math]

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
    [math]\ce{ Fe3O4_(s) {+~} 8{{nbsp}}HCl_(aq) -> FeCl2_(aq) {+~} 2{{nbsp}}FeCl3_(aq) {+~} 4{{nbsp}}H2O_(l) }[/math]
  2. Oxidation of iron(II) chloride with chlorine
    [math]\ce{ 2{{nbsp}}FeCl2_(aq) {+~} Cl2_(g) -> 2{{nbsp}}FeCl3_(aq) }[/math]
  3. Oxidation of iron(II) chloride with oxygen
    [math]\ce{ 4{{nbsp}}FeCl2_(aq) {+~} O2 {+~} 4{{nbsp}}HCl -> 4{{nbsp}}FeCl3_(aq) {+~} 2{{nbsp}}H2O_(l) }[/math]

Heating hydrated iron(III) chloride does not yield anhydrous ferric chloride. Instead, the solid decomposes into HCl and iron oxychloride. Hydrated iron(III) chloride can be converted to the anhydrous form by treatment with thionyl chloride.[12] Similarly, dehydration can be effected with trimethylsilyl chloride:[13]

[math]\ce{ FeCl3.6H2O + 12 Me3SiCl -> FeCl3 + 6 (Me3Si)2O + 12 HCl }[/math]


A brown, acidic solution of iron(III) chloride

When dissolved in water, iron(III) chloride give a strongly acidic solution.[14][9]

When heated with iron(III) oxide at 350 °C, iron(III) chloride gives iron oxychloride.[15]

[math]\ce{ FeCl3 + Fe2O3 -> 3FeOCl }[/math]

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

Redox reactions

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

[math]\ce{ FeCl3 + CuCl -> FeCl2 + CuCl2 }[/math]

In a comproportionation reaction, it reacts with iron to form iron(II) chloride:

[math]\ce{ 2 FeCl3 + Fe -> 3 FeCl2 }[/math]

A traditional synthesis of anhydrous ferrous chloride is the reduction of FeCl3 with chlorobenzene:[16]

[math]\ce{ 2 FeCl3 + C6H5Cl -> 2 FeCl2 + C6H4Cl2 + HCl }[/math]

With carboxylate anions

Oxalates react rapidly with aqueous iron(III) chloride to give [Fe(C
. 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.[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 FeCl
and sodium ethoxide:[19][20]

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

With organometallic compounds

Iron(III) chloride in ether solution oxidizes methyl lithium LiCH
to give first light greenish yellow lithium tetrachloroferrate(III) LiFeCl
solution and then, with further addition of methyl lithium, lithium tetrachloroferrate(II) Li

[math]\ce{ 2 FeCl3 + LiCH3 -> FeCl2 + LiFeCl4 + .CH3 }[/math]
[math]\ce{ LiFeCl4 + LiCH3 -> Li2FeCl4 + .CH3 }[/math]

The methyl radicals combine with themselves or react with other components to give mostly ethane C
and some methane CH



Iron(III) chloride is used in sewage treatment and drinking water production as a coagulant and flocculant.[23] In this application, FeCl
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.

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

It is also used as a leaching agent in chloride hydrometallurgy,[24] for example in the production of Si from FeSi (Silgrain process).[25]

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.[26]

[math]\ce{ FeCl3 + Cu -> FeCl2 + CuCl }[/math]
[math]\ce{ FeCl3 + CuCl -> FeCl2 + CuCl2 }[/math]

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.

[math]\ce{ H2C=CH2 + Cl2 -> ClCH2CH2Cl }[/math]

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.(citation?) It is less powerful than aluminium chloride, but in some cases this mildness leads to higher yields, for example in the alkylation of benzene:

Iron(III) chloride as a catalyst

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.

This test can be used to detect the presence of gamma-hydroxybutyric acid and gamma-butyrolactone,[28] 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.[29]
  • 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.[30]
  • Used in an animal thrombosis model.[31]
  • Used in energy storage systems.[32]
  • Historically it was used to make direct positive blueprints.[33][34]
  • A component of modified Carnoy's solution used for surgical treatment of keratocystic odontogenic tumor (KOT).


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]

FeCl3 is also produced as an atmospheric salt aerosol by reaction between iron-rich dust and hydrochloric acid from sea salt. This iron salt aerosol causes about 5% of naturally-occurring oxidisation of methane and is thought to have a range of cooling effects. [37]

See also


  1. An alternative GHS classification from the Japanese GHS Inter-ministerial Committee (2006)[7] notes the possibility of respiratory tract irritation from FeCl
    and differs slightly in other respects from the classification used here.


  1. 1.0 1.1 1.2 1.3 1.4 1.5 Haynes, William M., ed (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.69. ISBN 1439855110. 
  2. Haynes, William M., ed (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.133. ISBN 1439855110. 
  3. 3.0 3.1 "Structure refinement of an FeCl
    crystal using a thin plate sample". J. Appl. Crystallogr. 22 (2): 173. 1989. doi:10.1107/S0021889888013913.
  4. NIOSH Pocket Guide to Chemical Hazards. "#0346". National Institute for Occupational Safety and Health (NIOSH). 
  5. HSNO Chemical Classification Information Database, New Zealand Environmental Risk Management Authority,, retrieved 19 Sep 2010 
  6. Various suppliers, collated by the Baylor College of Dentistry, Texas A&M University. (accessed 2010-09-19)
  7. Template:GHS class JP
  8. Inorganic Chemistry. San Diego: Academic Press. 2001. ISBN 978-0-12-352651-9. 
  9. 9.0 9.1 9.2 Simon A. Cotton (2018). "Iron(III) chloride and its coordination chemistry". Journal of Coordination Chemistry 71 (21): 3415–3443. doi:10.1080/00958972.2018.1519188. 
  10. Lind, M. D. (1967). "Crystal Structure of Ferric Chloride Hexahydrate". The Journal of Chemical Physics 47 (3): 990–993. doi:10.1063/1.1712067. Bibcode1967JChPh..47..990L. 
  11. "Anhydrous Iron(III) Chloride". Inorganic Syntheses 3: 191–194. 1950. doi:10.1002/9780470132340.ch51. 
  12. Anhydrous Metal Chlorides. 28. 1990. pp. 321–323. doi:10.1002/9780470132593.ch80. ISBN 9780470132593. 
  13. "Solvated and Unsolvated Anhydrous Metal Chlorides from Metal Chloride Hydrates". Inorganic Syntheses 29: 108–111. 1992. doi:10.1002/9780470132609.ch26. 
  14. Housecroft, C. E.; Sharpe, A. G. (2012). Inorganic Chemistry (4th ed.). Prentice Hall. pp. 747. ISBN 978-0-273-74275-3. 
  15. "Layered Intercalation Compounds". Inorganic Syntheses. John Wiley & Sons, Inc.. 1984. pp. 86–89. doi:10.1002/9780470132531.ch17. ISBN 9780470132531. 
  16. P. Kovacic and N. O. Brace (1960). "Iron(II) Chloride". Inorganic Syntheses 6: 172–173. doi:10.1002/9780470132371.ch54. 
  17. "12.22.1 Synthesis". The Chemistry of Metal Alkoxides. Springer Science. 2002. pp. 481. ISBN 0306476576. 
  18. "3.2.10. Alkoxides of later 3d metals". Alkoxo and aryloxo derivatives of metals. San Diego: Academic Press. 2001. pp. 69. ISBN 9780121241407. OCLC 162129468. 
  19. "Fe
    —A New Structure Type of an Uncharged Iron(III) Oxide-Alkoxide Cluster". Eur. J. Inorg. Chem. 2001 (2): 367. 2001. doi:10.1002/1099-0682(200102)2001:2<367::AID-EJIC367>3.0.CO;2-V.
  20. "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. 2005. doi:10.1016/j.ica.2005.03.048. 
  21. "Über die Bildung von Lithiumtetrachloroferrat(II) Li
    bei der Umsetzung von Eisen(III)-chlorid mil Lithiummethyl (1:1) in ätherischer Lösung" (in de). Z. Anorg. Allg. Chem. 391 (3): 193–202. 1972. doi:10.1002/zaac.19723910302.
  22. "Über die Bildung von Lithiumtetrachloroferrat(II) Li
    bei der Umsetzung von Eisen(III)‐chlorid mil Lithiummethyl (1:1) in ätherischer Lösung" (in de). Z. Anorg. Allg. Chem. 391 (3): 193–202. 1972. doi:10.1002/zaac.19723910302.
  23. Water Treatment Chemicals. Akzo Nobel Base Chemicals. 2007. Retrieved 26 Oct 2007. 
  24. "A study on the acidified ferric chloride leaching of a complex (Cu–Ni–Co–Fe) matte". Separation and Purification Technology 51 (3): 332–337. 2006. doi:10.1016/j.seppur.2006.02.013. 
  25. "Validation of a compartmental population balance model of an industrial leaching process: The Silgrain process". Chem. Eng. Sci. 61 (1): 229–245. 2006. doi:10.1016/j.ces.2005.01.047. 
  26. Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. 1997. pp. 1084. ISBN 9780750633659. 
  27. Vogel's Textbook of Practical Organic Chemistry (5th ed.). New York: Longman/John Wiley & Sons. 1989. ISBN 9780582462366. 
  28. "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. 2006. PMID 17285863. 
  29. "A New Type of Organo-Iron Compound". Nature 168 (4285): 1040. 1951. doi:10.1038/1681039b0. Bibcode1951Natur.168.1039K. 
  30. "Mild and efficient reduction of azides to amines: synthesis of fused [2,1-b]quinazolines". Tetrahedron Lett. 43 (38): 6861–6863. 2002. doi:10.1016/S0040-4039(02)01454-5. 
  31. "Transendothelial migration of ferric ion in FeCl
    injured murine common carotid artery". Thromb. Res. 118 (2): 275–280. 2006. doi:10.1016/j.thromres.2005.09.004. PMID 16243382.
  32. Manohar, Aswin K.; Kim, Kyu Min; Plichta, Edward; Hendrickson, Mary; Rawlings, Sabrina; Narayanan, S. R. (28 October 2015). "A High Efficiency Iron-Chloride Redox Flow Battery for Large-Scale Energy Storage" (in en). Journal of the Electrochemical Society 163 (1): A5118. doi:10.1149/2.0161601jes. ISSN 1945-7111. 
  33. "Method of preparing paper" US patent Patent 241713, published 1881
  34. Modern Heliographic Processes. New York: D. Van Norstrand Company. 1888. pp. 65. 

Further reading

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