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 (anhydrous)
Iron(III) chloride (hydrate)
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 (aqueous solution)
  • 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
  • Cl[Fe](Cl)Cl
Properties
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 aqueous 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 (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)
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[4][5][Note 1]
GHS labelling:
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[7]
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 YN ?)
Infobox references

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 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]

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]

  • 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 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]

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

Reactions

Ferric chloride undergoes reactions according to two different chemical properties: It is a Lewis acid, and an oxidizing agent. In addition, it undergoes exchange reactions of its consitituent atoms.

Lewis-acid reactions

A brown, acidic solution of iron(III) chloride

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

Many oxo-anion chemicals react with ferric chloride to form simple ferric salts or complexes. Alkali metal alkoxides react to give the metal alkoxide complexes of varying complexity.[15] The compounds can be dimeric or trimeric.[16] In the solid phase a variety of multinuclear complexes have been described for the nominal stoichiometric reaction between FeCl3 and sodium ethoxide:[17][18]

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

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

The anhydrous salt forms 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 abstract a fourth chloride, give the salts of the yellow tetrachloroferrate ion ([FeCl4]), which easily dissolve in non-aqueous solvents. Combinations of FeCl3 with NaCl or KCl give Na[FeCl4] or K[FeCl4], respectively, rather than simply a binary mixture of the two chemicals.[19]

In addition to these simple stoichiometric reactions, the Lewis acidity of ferric chloride enables its use in a variety of acid-catalyzed reactions.

Redox reactions

Iron(III) chloride is a mild oxidizing agent. It undergoes simple one-electron reactions, such as oxidize 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:[20]

2 FeCl3 + C6H5Cl → 2 FeCl2 + C6H4Cl2 + HCl

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.

Attempts to perform a similar reaction using cyclopentadienyl magnesium bromide to form dimerized cyclopentadienyl instead gave ferrocene, an important reaction in the history of organometallic chemistry.[22][23]

Reaction with various neutral benzene derivatives can give chlorination of the benzene ring, via an apparent electrophilic aromatic substitution, or dimerization to form biphenyl compounds.[24]

Exchange reactions

When heated with iron(III) oxide at 350 °C, iron(III) chloride gives iron oxychloride.[25]

FeCl3 + Fe2O3 → 3FeOCl

Uses

Industrial

Iron(III) chloride is used in sewage treatment and drinking water production as a coagulant and flocculant.[26] 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

Iron(III) chloride is also used to remove soluble phosphate from wastewater. Upon addition, iron(III) phosphate forms. This salt is insoluble, and therefore easy to remove by later filtration and clarification stages of the purification process.[27]

It is also used as a leaching agent in chloride hydrometallurgy,[28] for example in the production of Si from FeSi (Silgrain process by Elkem).[29]

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).[30]

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

Catalysis

In the laboratory iron(III) chloride is commonly employed as a Lewis acid for catalysing reactions such as electrophilic aromatic substitutionchlorination of aromatic compounds and of aromatics. In this role, it is similar to aluminium chloride, and even sometimes used as mixtures with it.[31] It is somewhat milder, which can be an advantage in situations where it is important to prevent over-reaction, for example in the Friedel–Crafts reaction of benzene to give tert-butylbenzene:

Ferric chloride on silica gel is a reagent that has high reactivity towards several oxygen-containing functional groups. When the reagent is dry, its acidity and high affinity for water lead to dehydration and pinacol-type rearrangement reactions. When the reagent is moistened, it instead induces hydrolysis or epimerization reactions.[32]

Qualitative analysis

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.[33] 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[34],[35] which cause it to turn red-brown.

Medical applications

  • Used in veterinary practice to treat overcropping of an animal's claws, particularly when the overcropping results in bleeding
  • Used in an animal thrombosis model.[36]
  • A component of Carnoy's solution, a histological fixative with many applications

Metal etching

In addition to etching of copper on PC circuit boards, there are other applications involving etching of copper:

The comproportionation reaction finds use in etching of iron-containing materials:

Other etching applications include:

  • Used to strip aluminium coating from mirrors.[41]
  • Used to etch intricate medical devices.

Other uses

  • Used in anhydrous form as a drying reagent in certain reactions.
  • Used in wastewater treatment for odor control, by removal of hydrogen sulfide.[42]
  • Sometimes used in a technique of Raku ware firing,[43] the iron coloring a pottery piece shades of pink, brown, and orange.
  • Used in conjunction with NaI in acetonitrile to mildly reduce organic azides to primary amines.[44]
  • Used in an experimental energy storage systems.[45]
  • Historically it was used to make direct positive blueprints.[46][47]

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.[48][49]

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.[50]

The clouds of Venus are hypothesized to contain approximately 1% FeCl3 dissolved in sulfuric acid.[51][52]

See also

Notes

  1. An alternative GHS classification from the Japanese GHS Inter-ministerial Committee (2006)[6] notes the possibility of respiratory tract irritation from FeCl3 and differs slightly in other respects from the classification used here.

References

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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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