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)
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]
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[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 describes the inorganic compounds with the formula FeCl3(H2O)x. Also called ferric chloride, these compounds are available both in an anhydrous and hydrated forms. They are common source of iron in the +3 oxidation state. The hydrate and the anhydrous derivative have distinct properties.

Structure and properties

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

All forms of ferric chloride are paramagnetic, owing to the presence of five unpaired electrons residing in 3d orbitals. This electronic configuration places electrons in molecular orbitals that are antibonding with respect to ligands. Thus, iron(III) chlorides are labile, undergoing rapid ligand exchange in solution. In contrast to their kinetic lability, iron(III) chlorides are thermodynamically robust, as reflected by the vigorous methods applied to their synthesis.

Anhydrous

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. 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 aggressively forms hydrates upon exposure to water, reflecting its Lewis acidity. Four of these hydrates have been crystallized and examined by X-ray crystallography. They all feature trans-[FeCl2(H2O)4]+ cations, with either chloride or [FeCl4]− as the anions.[9]

Solution

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

Aqueous solutions of ferric chloride are characteristically yellow, in contrast to the pale pink solutions of [Fe(H2O)6]3+. Thus, the chloride ligand significantly influences the optical properties of the iron center. According to spectroscopic measurements, the main species in aqueous solutions of ferric chloride are the octahedral [FeCl2(H2O)4]+ (stereochemistry unspecified) and the tetrahedral [FeCl4]−.[9] The cationic aquo complex is strongly acidic:[11][9]

[FeCl2(H2O)4)]+ ⇌ [FeCl2(OH)(H2O)3] + H+

Anhydrous iron(III) chloride dissolves in diethyl ether and tetrahydrofuran forming 1:2 adducts of the formula FeCl3(ether)2. In these complexes, the iron is pentacoordinate.[12]

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:

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.[13] A similar but laboratory scale process also has been described.[14]

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.[15] Similarly, dehydration can be effected with trimethylsilyl chloride:[16]

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

Reactions

The reactivity of ferric chloride reveals two trends: It is a Lewis acid and an oxidizing agent.

Lewis-acid reactions

Reactions of iron(III) chlorides reflect the description of iron(III) as oxophilic and a hard Lewis acid. The rapidity of these reactions are consistent with the lability of ferric ion, reflecting its typical high-spin electronic configuration. Thus, oxalate salts react rapidly with aqueous iron(III) chloride to give [Fe(C2O4)3]3−, known as ferrioxalate. Other carboxylate sources, e.g., citrate and tartrate, bind as well to give carboxylate complexes. The affinity of iron(III) for oxygen ligands was the basis of qualitative tests for phenols. Although superceded by spectroscopic methods, the ferric chloride test is a traditional colorimetric test.[17] The affinity of iron(III) for phenols is exploited in the Trinder spot test.

Myriad other manifestation of the oxophiliicty of iron(III) chloride are available. When heated with iron(III) oxide at 350 °C it reactions to give iron oxychloride:[18]

FeCl3 + Fe2O3 → 3FeOCl

Alkali metal alkoxides react to give the iron(III) alkoxide complexes. These products have more complicated structures that the anhydrous iron(III) chloride.[19][20] 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 adduct with Lewis bases such as triphenylphosphine oxide; e.g., FeCl3(OP(C6H5)3)2. The related 1:2 complex , has been crystallized from ether solution.[12]

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.[21]

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.

Redox reactions

Iron(III) chloride is a mild oxidizing agent. It serves as one-electron oxidant illustrated by its reaction with copper(I) chloride to give copper(II) chloride and iron(II) chloride.

FeCl3 + CuCl → FeCl2 + CuCl2

In a comproportionation reaction, iron(III) chloride reacts with iron powder to form iron(II) chloride:[13]

2 FeCl3 + Fe → 3 FeCl2

A traditional synthesis of anhydrous ferrous chloride is the reduction of FeCl3 with chlorobenzene:[22]

2 FeCl3 + C6H5Cl → 2 FeCl2 + C6H4Cl2 + HCl

Organometallic chemistry

The interaction of anhydrous iron(III) chloride with organolithium and organomagnesium compounds has been examined often. These studies are enabled because of the solubility of FeCl3 in etherial solvents, which are compatible with the nucleophilic alkylating agents. Such studies may be relevant to the mechanism of FeCl3-catalyzed cross coupling reactions.[23] The isolation of organoiron(III) intermediates requires low temperature reactions, lest the [FeR4]- intermediates degrade. Using methylmagnesium bromide as the alkylation agent, salts of Fe(CH3)4]- have been isolated.[24] Illustrating the sensitivity of these reactions, methyl lithium LiCH3 reacts with iron(III) chloride to give lithium tetrachloroferrate(II) Li2[FeCl4]:[25]

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.[26] These diketonate complexes have the advantages that they do not form hydrates, unlike iron(III) chloride, and they are more soluble in relevant solvents.[23] Cyclopentadienyl magnesium bromide undergoes a complex reaction with iron(III) chloride, resulting in ferrocene:[27]

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.[28]

Uses

Water treatment

In the largest application iron(III) chloride is used in sewage treatment and drinking water production as a coagulant and flocculant.[29] In this application, an aqueous solution of FeCl3 is treated with base to form a floc of iron(III) hydroxide (Fe(OH)3), also formulated as FeO(OH) (ferrihydrite). This floc facilitates the separation of suspended materials, clarifying the water.[13]

Iron(III) chloride is also used to remove soluble phosphate from wastewater. Iron(III) phosphate is insoluble and thus precipitates as a solid.[30] One potential advantage to its use in water treatment, ferric ion oxidizes (deodorizes) hydrogen sulfide.[31]

Etching and metal cleaning

It is also used as a leaching agent in chloride hydrometallurgy,[32] for example in the production of Si from FeSi (Silgrain process by Elkem).[33]

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 copper(II) chloride is rinsed away, leaving a copper pattern. This chemistry is used in the production of printed circuit boards (PCB).[34]

Iron(III) chloride is used in many other hobbies involving metallic objects.[35][36][37][38][39]

Organic chemistry

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

In industry, iron(III) chloride is used as catalyst for the reaction of ethylene with chlorine, forming ethylene dichloride (1,2-dichloroethane):[40]

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.

Several reagents for organic synthesis have been developed based especially on anhydrous iron(III) chloride:

As a reagent in organic chemistry, iron(III) chloride has attracted interest for both its redox activity and its Lewis acidity. Furthermore, because they are inexpensive and relatively nontoxic, iron chlorides have been widely examined.[26] Illustrating it use as a Lewis acid, iron(III) chloride catalyses electrophilic aromatic substitution and chlorinations. In this role, its function is similar to that of aluminium chloride. In some cases, mixtures of the two are used.[45] Iron(III) chloride oxidizes naphthols to naphthoquinones:[26][46]

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.

Safety

Anhydrous iron(III) chloride is harmful, highly corrosive, and acidic.[26]

Natural occurrence

The natural counterpart of FeCl3 is the rare mineral molysite, usually related to volcanic and other-type fumaroles.[47][48]

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.[49]

The clouds of Venus are hypothesized to contain approximately 1% FeCl3 dissolved in sulfuric acid.[50][51]

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

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Further reading

Wikimedia Commons has media related to Iron(III) chloride.
  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. 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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