Iron(III) chloride

Iron(III) chloride is the inorganic compound with the formula FeCl3. 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 colour depends on the viewing angle: by reflected light the crystals appear dark green, but by transmitted light they appear purple-red.

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)
ECHA InfoCard 100.028.846 Edit this at Wikidata
EC Number
  • 231-729-4
RTECS number
  • LJ9100000
UN number
  • 1773 (anhydrous)
  • 2582 (aqueous solution)
  • InChI=1S/3ClH.Fe/h3*1H;/q;;;+3/p-3 checkY
  • InChI=1S/3ClH.Fe/h3*1H;/q;;;+3/p-3
  • Cl[Fe](Cl)Cl
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 aqueous 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 (anhydrous 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 labelling:
Corr. Met. 1; Skin Corr. 1C; Eye Dam. 1Acute Tox. 4 (oral)
H290, H302, H314
P234, P260, P264, P270, P273, P280, P301+P312, P301+P330+P331, P303+P361+P353, P304+P340, P305+P351+P338, P310, P321, P363, P390, P405, P406, P501
NFPA 704 (fire diamond)
Flash point Non-flammable
NIOSH (US health exposure limits):
REL (Recommended)
TWA 1 mg/m3[4]
Safety data sheet (SDS) ICSC 1499
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).
☒N verify (what is checkY☒N ?)
Iron(III) chloride hexahydrate
Iron(III) chloride hexahydrate

Structure and propertiesEdit


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 vapor consists of the dimer Fe2Cl6 (like aluminium chloride) which increasingly dissociates into the monomeric FeCl3 (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]

  • dihydrate: FeCl3·2H2O has the structural formula trans-[FeCl2(H2O)4][FeCl4].
  • FeCl3·2.5H2O has the structural formula cis-[FeCl2(H2O)4][FeCl4]·H2O.
  • FeCl3·3.5H2O has the structural formula cis-[FeCl2(H2O)4][FeCl4]·3H2O.
  • hexahydrate: FeCl3·6H2O has the structural formula trans-[FeCl2(H2O)4]Cl·2H2O.[10]

Aqueous solutionEdit

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]

2 Fe + 3 Cl2 → 2 FeCl3

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
    Fe3O4 + 8 HCl → FeCl2 + 2 FeCl3 + 4 H2O
  2. Oxidation of iron(II) chloride with chlorine
    2 FeCl2 + Cl2 → 2 FeCl3
  3. Oxidation of iron(II) chloride with oxygen and hydrochloric acid
    4 FeCl2 + O2 + 4 HCl → 4 FeCl3 + 2 H2O

Heating hydrated iron(III) chloride does not yield anhydrous ferric chloride. Instead, the solid decomposes into hydrochloric acid 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]

FeCl3·6H2O + 12 (CH3)3SiCl → FeCl3 + 6 ((CH3)3Si)2O + 12 HCl


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]

FeCl3 + Fe2O3 → 3FeOCl

The anhydrous salt is a moderately strong Lewis acid, forming adducts with Lewis bases such as triphenylphosphine oxide; e.g., FeCl3(OPPh3)2 where Ph is 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.

Redox reactionsEdit

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

FeCl3 + CuCl → FeCl2 + CuCl2

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

2 FeCl3 + Fe → 3 FeCl2

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

2 FeCl3 + C6H5Cl → 2 FeCl2 + C6H4Cl2 + HCl

With carboxylate anionsEdit

Oxalates react rapidly with aqueous iron(III) chloride to give [Fe(C2O4)3]3−. Other carboxylate salts form complexes; e.g., citrate and tartrate.

With alkali metal alkoxidesEdit

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 FeCl3 and sodium ethoxide:[19][20]

FeCl3 + 3 [CH3CH2O]Na+ → Fe(OCH2CH3)3 + 3 NaCl

With organometallic compoundsEdit

Iron(III) chloride in ether solution oxidizes methyl lithium LiCH3 to give first light greenish yellow lithium tetrachloroferrate(III) Li[FeCl4] solution and then, with further addition of methyl lithium, lithium tetrachloroferrate(II) Li2[FeCl4]:[21]

2 FeCl3 + LiCH3 → FeCl2 + Li[FeCl4] + •CH3
Li[FeCl4] + LiCH3 → Li2[FeCl4] + •CH3

The methyl radicals combine with themselves or react with other components to give mostly ethane C2H6 and some methane CH4.



Iron(III) chloride is used in sewage treatment and drinking water production as a coagulant and flocculant.[22] In this application, FeCl3 in slightly basic water reacts with the hydroxide ion (OH) to form a floc of iron(III) hydroxide (Fe(OH)3), also formulated as FeO(OH) (ferrihydrite), that can remove suspended materials.

[Fe(H2O)6]3+ + 4 OH → [Fe(OH)4(H2O)2] + 4 H2O → [FeO(OH)2(H2O)] + 6 H2O

It is also used as a leaching agent in chloride hydrometallurgy,[23] for example in the production of Si from FeSi (Silgrain process by Elkem).[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 (PCB).[25]

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 useEdit

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 needed] 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 neutralized 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 neutralized 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],[28] which cause it to turn red-brown.

Other usesEdit


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 occurrenceEdit

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 oxidization of methane and is thought to have a range of cooling effects.[37]

The clouds of Venus are hypothesized to contain approximately 1% FeCl3 dissolved in sulfuric acid.[38][39]

See alsoEdit


  1. ^ An alternative GHS classification from the Japanese GHS Inter-ministerial Committee (2006)[7] notes the possibility of respiratory tract irritation from FeCl3 and differs slightly in other respects from the classification used here.


  1. ^ a b c d e f Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.69. ISBN 1-4398-5511-0.
  2. ^ Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.133. ISBN 1-4398-5511-0.
  3. ^ a b Hashimoto S, Forster K, Moss SC (1989). "Structure refinement of an FeCl3 crystal using a thin plate sample". J. Appl. Crystallogr. 22 (2): 173–180. 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. ^ GHS classification – ID 831, Japanese GHS Inter-ministerial Committee, 2006, retrieved 19 Sep 2010
  8. ^ Holleman AF, Wiberg E (2001). Wiberg N (ed.). Inorganic Chemistry. San Diego: Academic Press. ISBN 978-0-12-352651-9.
  9. ^ a b c 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. S2CID 105925459.
  10. ^ Lind, M. D. (1967). "Crystal Structure of Ferric Chloride Hexahydrate". The Journal of Chemical Physics. 47 (3): 990–993. Bibcode:1967JChPh..47..990L. doi:10.1063/1.1712067.
  11. ^ Tarr BR, Booth HS, Dolance A (1950). Anhydrous Iron(III) Chloride. Inorganic Syntheses. Vol. 3. pp. 191–194. doi:10.1002/9780470132340.ch51.
  12. ^ Pray AR, Heitmiller RF, Strycker S, et al. (1990). "Anhydrous Metal Chlorides". Inorganic Syntheses. Vol. 28. pp. 321–323. doi:10.1002/9780470132593.ch80. ISBN 9780470132593.
  13. ^ Boudjouk P, So JH, Ackermann MN, et al. (1992). "Solvated and Unsolvated Anhydrous Metal Chlorides from Metal Chloride Hydrates". Inorganic Syntheses. Inorganic Syntheses. Vol. 29. pp. 108–111. doi:10.1002/9780470132609.ch26. ISBN 9780470132609.
  14. ^ Housecroft, C. E.; Sharpe, A. G. (2012). Inorganic Chemistry (4th ed.). Prentice Hall. p. 747. ISBN 978-0-273-74275-3.
  15. ^ Kikkawa S, Kanamaru F, Koizumi M, et al. (1984). "Layered Intercalation Compounds". In Holt SL Jr (ed.). Inorganic Syntheses. John Wiley & Sons, Inc. pp. 86–89. doi:10.1002/9780470132531.ch17. ISBN 9780470132531.
  16. ^ P. Kovacic and N. O. Brace (1960). "Iron(II) Chloride". Inorganic Syntheses. Inorganic Syntheses. Vol. 6. pp. 172–173. doi:10.1002/9780470132371.ch54. ISBN 9780470132371.
  17. ^ Turova NY, Turevskaya EP, Kessler VG, et al., eds. (2002). "12.22.1 Synthesis". The Chemistry of Metal Alkoxides. Springer Science. p. 481. ISBN 0306476576.
  18. ^ Bradley DC, Mehrotra RC, Rothwell I, et al. (2001). "3.2.10. Alkoxides of later 3d metals". Alkoxo and aryloxo derivatives of metals. San Diego: Academic Press. p. 69. ISBN 9780121241407. OCLC 162129468.
  19. ^ Michael V, Grätz F, Huch V (2001). "Fe9O3(OC2H5)21•C2H5OH—A New Structure Type of an Uncharged Iron(III) Oxide-Alkoxide Cluster". Eur. J. Inorg. Chem. 2001 (2): 367. doi:10.1002/1099-0682(200102)2001:2<367::AID-EJIC367>3.0.CO;2-V.
  20. ^ Seisenbaeva GA, Gohil S, Suslova EV, et al. (2005). "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. doi:10.1016/j.ica.2005.03.048.
  21. ^ Berthold HJ, Spiegl HJ (1972). "Über die Bildung von Lithiumtetrachloroferrat(II) Li2FeCl4 bei der Umsetzung von Eisen(III)-chlorid mil Lithiummethyl (1:1) in ätherischer Lösung". Z. Anorg. Allg. Chem. (in German). 391 (3): 193–202. doi:10.1002/zaac.19723910302.
  22. ^ Water Treatment Chemicals (PDF). Akzo Nobel Base Chemicals. 2007. Archived from the original (PDF) on 13 August 2010. Retrieved 26 Oct 2007.
  23. ^ Park KH, Mohapatra D, Reddy BR (2006). "A study on the acidified ferric chloride leaching of a complex (Cu–Ni–Co–Fe) matte". Separation and Purification Technology. 51 (3): 332–337. doi:10.1016/j.seppur.2006.02.013.
  24. ^ Dueñas Díez M, Fjeld M, Andersen E, et al. (2006). "Validation of a compartmental population balance model of an industrial leaching process: The Silgrain process". Chem. Eng. Sci. 61 (1): 229–245. doi:10.1016/j.ces.2005.01.047.
  25. ^ Greenwood NN, Earnshaw A (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. p. 1084. ISBN 9780750633659.
  26. ^ Furniss BS, Hannaford AJ, Smith PW, et al. (1989). Vogel's Textbook of Practical Organic Chemistry (5th ed.). New York: Longman/Wiley. ISBN 9780582462366.
  27. ^ "Gamma-Butyrolactone (GBL): All You Need to Know". Anbu Chem. Retrieved 2023-01-17.{{cite web}}: CS1 maint: url-status (link)
  28. ^ Zhang SY, Huang ZP (2006). "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. PMID 17285863.
  29. ^ Kealy TJ, Pauson PL (1951). "A New Type of Organo-Iron Compound". Nature. 168 (4285): 1040. Bibcode:1951Natur.168.1039K. doi:10.1038/1681039b0. S2CID 4181383.
  30. ^ Kamal A, Ramana K, Ankati H, et al. (2002). "Mild and efficient reduction of azides to amines: synthesis of fused [2,1-b]quinazolines". Tetrahedron Lett. 43 (38): 6861–6863. doi:10.1016/S0040-4039(02)01454-5.
  31. ^ Tseng M, Dozier A, Haribabu B, et al. (2006). "Transendothelial migration of ferric ion in FeCl3 injured murine common carotid artery". Thromb. Res. 118 (2): 275–280. 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". Journal of the Electrochemical Society. 163 (1): A5118. doi:10.1149/2.0161601jes. ISSN 1945-7111. S2CID 100823390.
  33. ^ US Patent 241713, Pellet H, "Method of preparing paper", published 1881 
  34. ^ Lietze E (1888). Modern Heliographic Processes. New York: D. Van Norstrand Company. pp. 65.
  35. ^ "Molysite".
  36. ^ "List of Minerals". March 21, 2011.
  37. ^ Oeste, Franz Dietrich; de Richter, Renaud; Ming, Tingzhen; Caillol, Sylvain (January 13, 2017). "Climate engineering by mimicking natural dust climate control: the iron salt aerosol method". Earth System Dynamics. 8 (1): 1–54. Bibcode:2017ESD.....8....1O. doi:10.5194/esd-8-1-2017 – via
  38. ^ Krasnopolsky, V. A.; Parshev, V. A. (1981). "Chemical composition of the atmosphere of Venus". Nature. 292 (5824): 610–613. Bibcode:1981Natur.292..610K. doi:10.1038/292610a0. S2CID 4369293.
  39. ^ Krasnopolsky, Vladimir A. (2006). "Chemical composition of Venus atmosphere and clouds: Some unsolved problems". Planetary and Space Science. 54 (13–14): 1352–1359. Bibcode:2006P&SS...54.1352K. doi:10.1016/j.pss.2006.04.019.

Further readingEdit

  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.