Iron(III) chloride

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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 Edit this at Wikidata
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 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
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)
Hazards[4][5][Note 1]
GHS labelling:
Corr. Met. 1; Skin Corr. 1C; Eye Dam. 1Acute Tox. 4 (oral)
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)
NFPA 704 four-colored diamondHealth 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
2
0
0
Flash point Non-flammable
NIOSH (US health exposure limits):
REL (Recommended)
 TWA 1 mg/m3[3]
Safety data sheet (SDS)
ICSC 1499
Related compounds
Other anions
Other cations
Related coagulants
Structure
Hexagonal, hR24
R3, No. 148[7]
a = 0.6065 nm, b = 0.6065 nm, c = 1.742 nm
α = 90°, β = 90°, γ = 120°
6
Octahedral
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 describes the inorganic compounds with the formula

Lewis acid, while all forms are mild oxidizing agents. It is used as a water cleaner and as an etchant
for metals.

Electronic and optical properties

Anhydrous iron(III) chloride evaporates at relatively mild temperatures to give the bitetrahedral dimer.

All forms of ferric chloride are

ferric nitrate, which contain [Fe(H2O)6]3+, are nearly colorless, whereas the chloride solutions are yellow. Thus, the chloride ligands significantly influence the optical properties of the iron center.[8][9]

Structure

Iron(III) chloride can exist as an anhydrous material and a series of hydrates, which results in distinct structures.

Anhydrous

The

octahedral Fe(III) centres interconnected by two-coordinate chloride ligands.[7][10]

Iron(III) chloride has a relatively low melting point and boils at around 315 °C. The vapor consists of the

monomeric FeCl3 (with D3h point group molecular symmetry) at higher temperatures, in competition with its reversible decomposition to give iron(II) chloride and chlorine gas.[11]

Hydrates

Ferric chloride form

trans or cis stereochemistry, reflecting the relative location of the chloride ligands on the octahedral Fe center. Four hydrates have been characterized by X-ray crystallography: the dihydrate FeCl3·2H2O, the disesquihydrate FeCl3·2.5H2O, the trisesquihydrate FeCl3·3.5H2O, and finally the hexahydrate FeCl3·6H2O. These species differ with respect to the stereochemistry of the octahedral iron cation, the identity of the anions, and the presence or absence of water of crystallization.[9] The structural formulas are [trans−FeCl2(H2O)4][FeCl4], [cis−FeCl2(H2O)4][FeCl4]·H2O, [cis−FeCl2(H2O)4][FeCl4]·H2O, and [trans−FeCl2(H2O)4]Cl·2H2O. The first three members of this series have the tetrahedral tetrachloroferrate ([FeCl4]−) anion.[12]

Solution

A brown, acidic solution of iron(III) chloride.

Like the solid hydrates, aqueous solutions of ferric chloride also consist of the octahedral [FeCl2(H2O)4]+ of unspecified stereochemistry.

molecular weight of 104, which exhibit the property of "aging", i.e., the structure change or evolve over the course of days.[13]
The polymeric species formed by the hydrolysis of ferric chlorides are key to the use of ferric chloride for water treatment.

In contrast to the complicated behavior of its aqueous solutions, solutions of iron(III) chloride in diethyl ether and tetrahydrofuran are well-behaved. Both ethers form 1:2 adducts of the general formula FeCl3(ether)2. In these complexes, the iron is pentacoordinate.[14]

Preparation

Several hundred thousand kilograms of anhydrous iron(III) chloride are produced annually. The principal method, called direct chlorination, uses scrap iron as a precursor:[10]

2 Fe + 3 Cl2 → 2 FeCl3

The reaction is conducted at several hundred degrees such that the product is gaseous. Using excess chlorine guarantees that the intermediate ferrous chloride is converted to the ferric state.[10] A similar but laboratory-scale process also has been described.[15][16]

Aqueous solutions of iron(III) chloride are also produced industrially from a number of iron precursors, including iron oxides:

Fe2O3 + 12 HCl + 9 H2O → FeCl3(H2O)6

In complementary route, iron metal can be oxidized by hydrochloric acid followed by chlorination:[10]

Fe + 2 HCl → FeCl2 + H2
FeCl2 + 0.5 Cl2 + 6 H2O → FeCl3(H2O)6

A number of variables apply to these processes, including the oxidation of iron by ferric chloride and the hydration of intermediates.[10] Hydrates of iron(III) chloride do not readily yield anhydrous ferric chloride. Attempted thermal dehydration yields hydrochloric acid and iron oxychloride. In the laboratory, hydrated iron(III) chloride can be converted to the anhydrous form by treatment with thionyl chloride[17] or trimethylsilyl chloride:[18]

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

Reactions

Being

labile, meaning that its Cl- and H2O ligands exchange rapidly with free chloride and water.[9][19]
In contrast to their kinetic lability, iron(III) chlorides are thermodynamically robust, as reflected by the vigorous methods applied to their synthesis, as described above.

Anhydrous FeCl3

Aside from lability, which applies to anhydrous and hydrated forms, the reactivity of anhydrous ferric chloride reveals two trends: It is a

Lewis acid and an oxidizing agent.[20]

Reactions of anhydrous iron(III) chloride reflect its description as both

oxophilic and a hard Lewis acid. Myriad manifestations of the oxophiliicty of iron(III) chloride are available. When heated with iron(III) oxide at 350 °C it reacts to give iron oxychloride:[21]

FeCl3 + Fe2O3 → 3FeOCl

Alkali metal

alkoxide complexes. These products have more complicated structures than anhydrous iron(III) chloride.[22][23] In the solid phase a variety of multinuclear complexes have been described for the nominal stoichiometric reaction between FeCl3 and sodium ethoxide
:

FeCl3 + 3 CH3CH2ONa → "Fe(OCH2CH3)3" + 3 NaCl

Iron(III) chloride forms a 1:2

Lewis bases such as triphenylphosphine oxide; e.g., FeCl3(OP(C6H5)3)2. The related 1:2 complex FeCl3(OEt2)2, where Et = C2H5), has been crystallized from ether solution.[14]

Iron(III) chloride also reacts with tetraethylammonium chloride to give the yellow salt of the tetrachloroferrate ion ((Et4N)[FeCl4]). Similarly, combining FeCl3 with NaCl and KCl gives Na[FeCl4] and K[FeCl4], respectively.[24]

In addition to these simple

stoichiometric reactions, the Lewis acidity of ferric chloride enables its use in a variety of acid-catalyzed reactions as described below in the section on organic chemistry.[10]

In terms of its being an oxidant, iron(III) chloride oxidizes iron powder to form iron(II) chloride via a comproportionation reaction:[10]

2 FeCl3 + Fe → 3 FeCl2

A traditional synthesis of anhydrous

FeCl3 with chlorobenzene:[25]

2 FeCl3 + C6H5Cl → 2 FeCl2 + C6H4Cl2 + HCl

iron(III) chloride releases chlorine gas when heated above 160 °C, generating

ferrous chloride:[16]

2FeCl3 → 2FeCl2 + Cl2

To suppress this reaction, the preparation of iron(III) chloride requires an excess of chlorinating agent, as discussed above.[16][10]

Hydrated FeCl3

Unlike the anhydrous material, hydrated ferric chloride is not a particularly strong Lewis acid since water ligands have quenched the Lewis acidity by binding to Fe(III).

Like the anhydrous material, hydrated ferric chloride is oxophilic. For example,

colorimetric test.[26] The affinity of iron(III) for phenols is exploited in the Trinder spot test.[27]

Aqueous iron(III) chloride serves as a one-electron oxidant illustrated by its reaction with copper(I) chloride to give copper(II) chloride and iron(II) chloride.

FeCl3 + CuCl → FeCl2 + CuCl2

This fundamental reaction is relevant to the use of ferric chloride solutions in etching copper.

Organometallic chemistry

The interaction of anhydrous iron(III) chloride with

methyl lithium LiCH3 reacts with iron(III) chloride to give lithium tetrachloroferrate(II) Li2[FeCl4]:[30]

2 FeCl3 + LiCH3 → FeCl2 + Li[FeCl4] + 0.5 CH3CH3
Li[FeCl4] + LiCH3 → Li2[FeCl4] + 0.5 CH3CH3

To a significant extent,

iron(III) acetylacetonate and related beta-diketonate complexes are more widely used than FeCl3 as ether-soluble sources of ferric ion.[20] These diketonate complexes have the advantages that they do not form hydrates, unlike iron(III) chloride, and they are more soluble in relevant solvents.[28]
Cyclopentadienyl magnesium bromide undergoes a complex reaction with iron(III) chloride, resulting in ferrocene:[31]

3 C5H5MgBr + FeCl3 → Fe(C5H5)2 + 1/n (C5H5)n + 3 MgBrCl

This conversion, although not of practical value, was important in the history of organometallic chemistry where ferrocene is emblematic of the field.[32]

Uses

Water treatment

In the largest application iron(III) chloride is in

iron(III) hydroxide (Fe(OH)3), also formulated as FeO(OH) (ferrihydrite). This floc facilitates the separation of suspended materials, clarifying the water.[10]

Iron(III) chloride is also used to remove soluble

insoluble and thus precipitates as a solid.[34] One potential advantage to its use in water treatment, ferric ion oxidizes (deodorizes) hydrogen sulfide.[35]

Etching and metal cleaning

It is also used as a leaching agent in chloride hydrometallurgy,[36] for example in the production of Si from FeSi (Silgrain process by Elkem).[37]

In another commercial application, a solution of iron(III) chloride is useful for etching copper according to the following equation:

2 FeCl3 + Cu → 2 FeCl2 + CuCl2

The soluble

printed circuit boards (PCB).[19]

Iron(III) chloride is used in many other hobbies involving metallic objects.[38][39][40][41][42]

Organic chemistry

Structure of FeCl3(diethylether)2. Color code: Cl=green,Fe = blue, O = red.

In industry, iron(III) chloride is used as a catalyst for the reaction of

1,2-dichloroethane):[43]

H2C=CH2 + Cl2 → ClCH2CH2Cl

Ethylene dichloride is a commodity chemical, which is mainly used for the industrial production of vinyl chloride, the monomer for making PVC.[44]

Illustrating it use as a

chlorinations. In this role, its function is similar to that of aluminium chloride. In some cases, mixtures of the two are used.[45]

Organic synthesis research

Although iron(III) chlorides are seldom used in practical organic synthesis, they have received considerable attention as reagents because they are inexpensive, earth abundant, and relatively nontoxic. Many experiments probe both its redox activity and its Lewis acidity.[20] For example, iron(III) chloride oxidizes naphthols to naphthoquinones:[20][46] 3-Alkylthiophenes are polymerized to polythiophenes upon treatment with ferric chloride.[47] Iron(III) chloride has been shown to promote C-C coupling reaction.[48]

Several reagents have been developed based on

alumina, ferric chloride has been shown to accelerate ene reactions.[50]

When pretreated with sodium hydride, iron(III) chloride gives a hydride reducing agent that convert alkenes and ketones into alkanes and alcohols, respectively.[51]

Histology

Iron(III) chloride is a component of useful stains, such as Carnoy's solution, a histological fixative with many applications. Also, it is used to prepare Verhoeff's stain.[52]

Natural occurrence

Like many metal halides, FeCl3 naturally occurs as a trace mineral. The rare mineral molysite is usually associated with volcanoes and fumaroles.[53][54]

FeCl3-based aerosol are produced by a reaction between iron-rich dust and hydrochloric acid from sea salt. This iron salt aerosol causes about 1-5% of naturally-occurring oxidization of methane and is thought to have a range of cooling effects; thus, it has been proposed as a catalyst for Atmospheric Methane Removal.[55]

The clouds of Venus are hypothesized to contain approximately 1% FeCl3 dissolved in sulfuric acid.[56][57]

Safety

Iron(III) chlorides are widely used in

corrosive, and must be handled using proper protective equipment.[20]

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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  4. ^ HSNO Chemical Classification Information Database, New Zealand Environmental Risk Management Authority, retrieved 19 Sep 2010
  5. ^ Various suppliers, collated by the Baylor College of Dentistry, Texas A&M University. (accessed 2010-09-19)
  6. ^ GHS classification – ID 831, Japanese GHS Inter-ministerial Committee, 2006, retrieved 19 Sep 2010
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  16. ^ a b c H. Lux (1963). "Iron (III) Chloride". In G. Brauer (ed.). Handbook of Preparative Inorganic Chemistry, 2nd Ed. Vol. 2. NY, NY: Academic Press. p. 1492.
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  34. ^ "Phosphorus Treatment and Removal Technologies" (PDF). Minnesota Pollution Control Agency. June 2006.
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  38. ^ John David Graham. "Safer Printmaking—Intaglio". University of Saskatchewan. Retrieved 5 February 2024.
  39. ^ Harris P, Hartman R, Hartman J (November 1, 2002). "Etching Iron Meteorites". Meteorite Times. Retrieved October 14, 2016.
  40. ^ Mike Lockwood, Carl Zambuto. "A message about mirror coating and recoating". Lockwood Custom Optics, Inc. Lockwood Custom Optics. Retrieved 5 February 2024.
  41. ^ CoinValueLookup. "Buffalo Nickel No Date Value: How Much Is It Worth Today?". CoinValueLookup. CoinValueLookup. Retrieved 5 February 2024.
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Further reading

Wikimedia Commons has media related to Iron(III) chloride.
  1. Lide DR, ed. (1990).
    ISBN 9780849304712
    .
  2. Stecher PG, Finkel MJ, Siegmund OH, eds. (1960). The Merck Index of Chemicals and Drugs (7th ed.). Rahway, New Jersey, US: 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. 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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