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Iron(III) chloride

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Iron(III) chloride
Iron(III) chloride hexahydrate.jpg
Iron(III) chloride (hydrate)
Iron(III) chloride anhydrate.jpg
Iron(III) chloride (anhydrous)
Iron-trichloride-sheet-3D-polyhedra.png
Iron-trichloride-sheets-stacking-3D-polyhedra.png
Names
IUPAC names
Iron(III) chloride
Iron trichloride
Other names
  • Ferric chloride
  • Molysite
  • Flores martis
Identifiers
  • 7705-08-0 checkY
  • 10025-77-1 (hexahydrate) checkY
  • 54862-84-9 (dihydrate) checkY
  • 64333-00-2 (3.5hydrate)
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.028.846 Edit this at Wikidata
EC Number
  • 231-729-4
RTECS number
  • LJ9100000
UNII
UN number
  • 1773 (anhydrous)
  • 2582 (aq. soln.)
  • InChI=1S/3ClH.Fe/h3*1H;/q;;;+3/p-3 checkY
    Key: RBTARNINKXHZNM-UHFFFAOYSA-K checkY
  • InChI=1S/3ClH.Fe/h3*1H;/q;;;+3/p-3
    Key: RBTARNINKXHZNM-DFZHHIFOAF
  • Key: RBTARNINKXHZNM-UHFFFAOYSA-K
  • Cl[Fe](Cl)Cl
Properties
Chemical formula
 FeCl3
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
Density
  • 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
  • Acetone
  • Methanol
  • Ethanol
  • Diethyl ether[1]
  •  
  • 630 g/L (18 °C)
  • Highly soluble
  • 830 g/L
  • Highly soluble
Magnetic susceptibility (χ)
 +13,450·10−6 cm3/mol[2]
Viscosity 12 cP (40% solution)
Structure
Crystal structure
 Hexagonal, hR24
Space group
 R3, No. 148[3]
a = 0.6065 nm, b = 0.6065 nm, c = 1.742 nm
α = 90°, β = 90°, γ = 120°
6
Octahedral
Hazards[5][6][Note 1]
Safety data sheet ICSC 1499
GHS labelling:
Corr. Met. 1; Skin Corr. 1C; Eye Dam. 1Acute Tox. 4 (oral)
Signal word
 Danger
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)
2
0
0
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).
☒N  (what is checkY☒N ?)
Infobox references

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 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) centers 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 (cf. aluminum 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]

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[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]

Preparation[]

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

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
  2. Oxidation of iron(II) chloride with chlorine
  3. Oxidation of iron(II) chloride with oxygen

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]

Reactions[]

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]

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 reactions[]

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

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

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

With carboxylate anions[]

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 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 FeCl3 and sodium ethoxide:[19][20]

With organometallic compounds[]

Iron(III) chloride in ether solution oxidizes methyl lithium LiCH3 to give first light greenish yellow lithium tetrachloroferrate(III) LiFeCl4 solution and then, with further addition of methyl lithium, lithium tetrachloroferrate(II) Li2FeCl4:[21][22]

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

Uses[]

Industrial[]

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

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]

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.

Laboratory use[]

In the laboratory iron(III) chloride is commonly employed as a Lewis acid for catalyzing reactions such as chlorination of aromatic compounds and Friedel–Crafts reaction of aromatics.[citation needed] It is less powerful than aluminum 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 neutralized 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 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,[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 synthesized 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 aluminum 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).

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.

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

See also[]

Notes[]

  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.

References[]

  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. doi:10.1107/S0021889888013913. templatestyles stripmarker in |title= at position 28 (help)
  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. 3. pp. 191–194. doi:10.1002/9780470132340.ch51.
  12. ^ Pray AR, Heitmiller RF, Strycker S, et al. (1990). "Anhydrous Metal Chlorides". Inorganic Syntheses. 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. 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. 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. templatestyles stripmarker in |title= at position 1 (help)
  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. templatestyles stripmarker in |title= at position 51 (help)
  22. ^ 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. templatestyles stripmarker in |title= at position 51 (help)
  23. ^ Water Treatment Chemicals (PDF). Akzo Nobel Base Chemicals. 2007. Archived from the original (PDF) on 13 August 2010. Retrieved 26 Oct 2007.
  24. ^ 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.
  25. ^ 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.
  26. ^ Greenwood NN, Earnshaw A (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. p. 1084. ISBN 9780750633659.
  27. ^ Furniss BS, Hannaford AJ, Smith PW, et al. (1989). Vogel's Textbook of Practical Organic Chemistry (5th ed.). New York: Longman/Wiley. ISBN 9780582462366.
  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. templatestyles stripmarker in |title= at position 45 (help)
  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.
  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". www.mindat.org.
  36. ^ "List of Minerals". www.ima-mineralogy.org. 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 esd.copernicus.org.

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.

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