Open main menu

Iron(III) chloride (FeCl
3
), also called ferric chloride, is an industrial scale commodity chemical compound with iron in the +3 oxidation state. The compound also exist as a hexahydrate with the formula trans-[Fe(H
2
O)
4
Cl
2
]Cl · 2H2O normally written as FeCl
3
 · 6H
2
O
. 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. The hexahydrate has a melting point of 37 °C and appears orange-brown in color. In nature, iron(III) chloride is known as the mineral molysite, but it is rare and mainly found from some fumaroles. It is however an industrial scale commodity.

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
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.028.846
EC Number 231-729-4
RTECS number LJ9100000
UNII
UN number
  • 1773 (anhydrous)
  • 2582 (aq. soln.)
Properties
FeCl
3
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
  •  
  • 630 g/L (18 °C)
  • Highly soluble
  • 830 g/L
  • Highly soluble
+13,450·10−6 cm3/mol[2]
Viscosity 12 cP (40% solution)
Structure
Hexagonal, hR24
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 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
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
0
2
0
Flash point Non-flammable
US health exposure limits (NIOSH):
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 verify (what is ☑Y☒N ?)
Infobox references

Iron(III) chloride dissolves in water, but undergoes partial hydrolysis in an exothermic reaction, and result in a strongly acidic solution. 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 deliquescent; the partial hydrolysis also occurs as it absorbs water from the air, liberating hydrogen chloride that forms mists in moist air. It is a fairly strong Lewis acid, and it is used as a catalyst in organic synthesis.

Contents

Structure and propertiesEdit

AnhydrousEdit

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
2
Cl
6
(cf. aluminium chloride) which increasingly dissociates into the monomeric FeCl
3
(with D3h point group molecular symmetry) at higher temperature, in competition with its reversible decomposition to give iron(II) chloride and chlorine gas.[8]

HexahydrateEdit

The hexahydrate is a orange-brown solid[9] usually given as the simplified formula FeCl
3
 · 6H2O. The compound contains a cationic coordination complex, and is more properly written as trans-[Fe(H
2
O)
4
Cl
2
]Cl · 2H2O with the systematic name tetraaquadichloroiron(III) chloride dihydrate. The two H
2
O molecules are embedded within the monoclinic crystal structure.[3][10]

PreparationEdit

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

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.[12] Similarly, dehydration can be effected with trimethylsilyl chloride:[13]

 

ReactionsEdit

 
A brown, acidic solution of iron(III) chloride

Iron(III) chloride undergoes hydrolysis to give a strongly acidic solution.[14]

When heated with iron(III) oxide at 350 °C, iron(III) chloride gives iron oxychloride, a layered solid and intercalation host.[15]

 

The anhydrous salt is a moderately strong Lewis acid, forming adducts with Lewis bases such as triphenylphosphine oxide; e.g., FeCl
3
(OPPh
3
)
2
where Ph is phenyl. It also reacts with other chloride salts to give the yellow tetrahedral [FeCl
4
]
ion. Salts of [FeCl
4
]
in hydrochloric acid can be extracted into diethyl ether.

OxidationEdit

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

 

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

 

Reducing agents such as hydrazine convert iron(III) chloride to complexes of iron(II).

DecompositionEdit

FeCl3 can be decomposed by electrolysis[16][17]:

 
 
Granulated iron(III) chloride hexahydrate

With carboxylate anionsEdit

Oxalates react rapidly with aqueous iron(III) chloride to give [Fe(C
2
O
4
)
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.[18] The compounds can be dimeric or trimeric.[19] In the solid phase a variety of multinuclear complexes have been described for the nominal stoichiometric reaction between FeCl
3
and sodium ethoxide:[20][21]

 

With organometallic compoundsEdit

Iron(III) chloride in ether solution oxidizes methyl lithium LiCH
3
to give first light greenish yellow lithium tetrachloroferrate(III) LiFeCl
4
solution and then, with further addition of methyl lithium, lithium tetrachloroferrate(II) Li
2
FeCl
4
:[22][23]

 
 

The methyl radicals combine with themselves or react with other components to give mostly ethane C
2
H
6
and some methane CH
4
.

UsesEdit

IndustrialEdit

Iron(III) chloride is used in sewage treatment and drinking water production.[24] In this application, FeCl
3
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,[25] for example in the production of Si from FeSi (Silgrain process).[26]

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.[27]

 
 

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 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 neutralised with sodium hydroxide until a slight precipitate of FeO(OH) is formed.[28] 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,[29] which cause it to turn red-brown.

Other usesEdit

  • 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 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).
  • 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.[30]
  • 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.[31]
  • Used in an animal thrombosis model.[32]
  • Used in energy storage systems.
  • 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).

SafetyEdit

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 alsoEdit

NotesEdit

  1. ^ An alternative GHS classification from the Japanese GHS Inter-ministerial Committee (2006)[7] notes the possibility of respiratory tract irritation from FeCl
    3
    and differs slightly in other respects from the classification used here.

ReferencesEdit

  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 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. ^ a b c Hashimoto S, Forster K, Moss SC (1989). "Structure refinement of an FeCl
    3
    crystal using a thin plate sample". J. Appl. Crystallogr. 22 (2): 173. 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. ^ Housecroft, C. E.; Sharpe, A. G. (2012). Inorganic Chemistry (4th ed.). Prentice Hall. p. 747. ISBN 978-0-273-74275-3.
  10. ^ Lind MD (1967). "Crystal Structure of Ferric Chloride Hexahydrate". J. Chem. Phys. 47 (3): 990–993. doi:10.1063/1.1712067.
  11. ^ Tarr BR, Booth HS, Dolance A (1950). "Anhydrous Iron(III) Chloride". In Audrieth LF (ed.). Inorganic Syntheses. 3. McGraw-Hill Book Company, Inc. pp. 191–194. doi:10.1002/9780470132340.ch51. ISBN 9780470132340.
  12. ^ Pray AR, Heitmiller RF, Strycker S, et al. (1990). "Anhydrous Metal Chlorides". In Angelici RJ (ed.). 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". In Grimes RN (ed.). 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. ^ Kanungo SB, Mishra SK (1996). "Thermal dehydration and decomposition of FeCl
    3
    · xH
    2
    O". J. Therm. Anal. Calorim. 46 (5): 1487–1500. doi:10.1007/BF01979262. ISSN 0022-5215.
  17. ^ "Iron trichloride". Pubchem. Retrieved 12 May 2018.
  18. ^ 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.
  19. ^ 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.
  20. ^ Michael V, Grätz F, Huch V (2001). "Fe
    9
    O
    3
    (OC
    2
    H
    5
    )
    21
    ·C
    2
    H
    5
    OH—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.
  21. ^ 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". Inorganica Chim. Acta. 358 (12): 3506. doi:10.1016/j.ica.2005.03.048.
  22. ^ Berthold HJ, Spiegl HJ (1972). "Über die Bildung von Lithiumtetrachloroferrat(II) Li
    2
    FeCl
    4
    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.
  23. ^ Berthold HJ, Spiegl HJ (1972). "Über die Bildung von Lithiumtetrachloroferrat(II) Li
    2
    FeCl
    4
    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.
  24. ^ Water Treatment Chemicals (PDF). Akzo Nobel Base Chemicals. 2007. Retrieved 26 Oct 2007.
  25. ^ 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. doi:10.1016/j.seppur.2006.02.013.
  26. ^ 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.
  27. ^ Greenwood NN, Earnshaw A (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. p. 1084. ISBN 9780750633659.
  28. ^ Furniss BS, Hannaford AJ, Smith PW, et al. (1989). Vogel's Textbook of Practical Organic Chemistry (5th ed.). New York: Longman/Wiley. ISBN 9780582462366.
  29. ^ 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.
  30. ^ Kealy TJ, Pauson PL (1951). "A New Type of Organo-Iron Compound". Nature. 168 (4285): 1040. doi:10.1038/1681039b0.
  31. ^ 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): 6961. doi:10.1016/S0040-4039(02)01454-5.
  32. ^ Tseng M, Dozier A, Haribabu B, et al. (2006). "Transendothelial migration of ferric ion in FeCl
    3
    injured murine common carotid artery". Thromb. Res. 118 (2): 275–280. doi:10.1016/j.thromres.2005.09.004. PMID 16243382.
  33. ^ ‹See Tfd›US Patent 241713, ‹See Tfd›Pellet H, "Method of preparing paper", published 1881 
  34. ^ Lietze E (1888). Modern Heliographic Processes. New York: D. Van Norstrand Company. p. 65.

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. p. 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.