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| IUPAC names Iron(III) chloride Iron trichloride | |||
Other names
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3D model (JSmol) | |||
| ChEBI | |||
| ChemSpider |
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| ECHA InfoCard | 100.028.846 | ||
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| Properties | |||
| FeCl3 | |||
| Molar mass |
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| Appearance | Green-black by reflected light; purple-red by transmitted light; yellow solid as hexahydrate; brown as aqueous solution | ||
| Odor | SlightHCl | ||
| Density |
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| Melting point | 307.6 °C (585.7 °F; 580.8 K) (anhydrous) 37 °C (99 °F; 310 K) (hexahydrate)[1] | ||
| Boiling point |
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| 912 g/L (anhydrous or hexahydrate, 25 °C)[1] | |||
| Solubility in |
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| +13,450·10−6 cm3/mol[2] | |||
| Viscosity | 12 cP (40% solution) | ||
| Hazards[4][5][Note 1] | |||
| GHS labelling: | |||
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| 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) | |||
| Flash point | Non-flammable | ||
| NIOSH (US health exposure limits): | |||
REL (Recommended) | TWA 1 mg/m3[3] | ||
| Safety data sheet (SDS) | ICSC1499 | ||
| Related compounds | |||
Otheranions | |||
Othercations | |||
Relatedcoagulants | |||
| 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° | |||
Formula units (Z) | 6 | ||
| Octahedral | |||
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |||
Iron(III) chloride describes the inorganic compounds with the formulaFeCl3(H2O)x. Also calledferric chloride, these compounds are some of the most important and commonplace compounds of iron. They are available both in anhydrous and in hydrated forms, which are bothhygroscopic. They feature iron in its +3oxidation state. The anhydrous derivative is aLewis acid, while all forms are mild oxidizing agents. It is used as awater cleaner and as anetchant for metals.
All forms of ferric chloride areparamagnetic, owing to the presence of unpaired electrons residing in 3d orbitals. Although Fe(III) chloride can be octahedral or tetrahedral (or both, see structure section), all of these forms have five unpaired electrons, one perd-orbital. Thehigh spin d5 electronic configuration requires that d-d electronic transitions arespin forbidden, in addition to violating theLaporte rule. This double forbidden-ness results in its solutions being only pale colored. Or, stated more technically, the optical transitions are non-intense. Aqueousferric sulfate andferric 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]
Iron(III) chloride can exist as an anhydrous material and a series of hydrates, which results in distinct structures.
Theanhydrous compound is a hygroscopic 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 bytransmitted light, they appear purple-red. Anhydrous iron(III) chloride has theBiI3 structure, withoctahedral Fe(III) centres interconnected by two-coordinate chlorideligands.[7][10]
Iron(III) chloride has a relatively low melting point, and boils at around 315 °C. The vapor consists of thedimerFe2Cl6, much likealuminium chloride. This dimer dissociates into themonomericFeCl3 (with D3hpoint groupmolecular symmetry) at higher temperatures, in competition with its reversible decomposition to giveiron(II) chloride andchlorine gas.[11]
Ferric chloride formhydrates upon exposure to water, reflecting its Lewis acidity. All hydrates exhibitdeliquescence, meaning that they become liquid by absorbing moisture from the air. Hydration invariably gives derivatives ofaquo complexes with the formula[FeCl2(H2O)4]+. This cation can adopt eithertrans orcisstereochemistry, reflecting the relative location of the chlorideligands on theoctahedral Fe center. Four hydrates have been characterized byX-ray crystallography: the dihydrateFeCl3·2H2O, the disesquihydrateFeCl3·2.5H2O, the trisesquihydrateFeCl3·3.5H2O, and finally the hexahydrateFeCl3·6H2O. These species differ with respect to the stereochemistry of the octahedral iron cation, the identity of the anions, and the presence or absence ofwater 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 tetrahedraltetrachloroferrate ([FeCl4]−) anion.[12]
_chloride_2.JPG)
Like the solid hydrates, aqueous solutions of ferric chloride also consist of the octahedral[FeCl2(H2O)4]+ of unspecified stereochemistry.[9] Detailed speciation of aqueous solutions of ferric chloride is challenging because the individual components do not have distinctive spectroscopic signatures. Iron(III) complexes, with a high spin d5 configuration, are kinetically labile, which means that ligands rapidly dissociate and reassociate. A further complication is that these solutions are strongly acidic, as expected foraquo complexes of a tricationic metal. Iron aquo complexes are prone toolation, the formation ofpolymericoxo derivatives. Dilute solutions of ferric chloride produce soluble nanoparticles withmolecular 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 indiethyl ether andtetrahydrofuran are well-behaved. Bothethers form 1:2adducts of the general formula FeCl3(ether)2. In these complexes, the iron is pentacoordinate.[14]
Several hundred tons of anhydrous iron(III) chloride are produced annually. The principal method, calleddirect chlorination, uses scrap iron as a precursor:[10]
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:
In complementary route, iron metal can be oxidized byhydrochloric acid followed by chlorination:[10]
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 andiron oxychloride. In the laboratory, hydrated iron(III) chloride can be converted to the anhydrous form by treatment withthionyl chloride[17] ortrimethylsilyl chloride:[18]
Beinghigh spin d5 electronic configuration iron(III) chlorides arelabile, 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.
Aside from lability, which applies to anhydrous and hydrated forms, the reactivity of anhydrous ferric chloride reveals two trends: It is aLewis acid and anoxidizing agent.[20]
Reactions of anhydrous iron(III) chloride reflect its description as bothoxophilic and ahard Lewis acid. Myriad manifestations of the oxophiliicty of iron(III) chloride are available. When heated withiron(III) oxide at 350 °C it reacts to giveiron oxychloride:[21]
Alkali metalalkoxides react to give the iron(III)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 betweenFeCl3 andsodium ethoxide:
Iron(III) chloride forms a 1:2adduct withLewis bases such astriphenylphosphine oxide; e.g.,FeCl3(OP(C6H5)3)2. The related 1:2 complexFeCl3(OEt2)2, where Et = C2H5), has been crystallized from ether solution.[14]
Iron(III) chloride also reacts withtetraethylammonium chloride to give the yellow salt of thetetrachloroferrate ion ((Et4N)[FeCl4]). Similarly, combining FeCl3 with NaCl and KCl givesNa[FeCl4] andK[FeCl4], respectively.[24]
In addition to these simplestoichiometric 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 acomproportionation reaction:[10]
A traditional synthesis of anhydrousferrous chloride is the reduction ofFeCl3 withchlorobenzene:[25]
iron(III) chloride releases chlorine gas when heated above 160 °C, generatingferrous chloride:[16]
To suppress this reaction, the preparation of iron(III) chloride requires an excess of chlorinating agent, as discussed above.[16][10]
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). Instead, it is a Brønsted-Lowry acid, as the hydrogen atoms on the water ligands become more acidic when the water ligands bond to Fe(III).
Like the anhydrous material, hydrated ferric chloride is oxophilic. For example,oxalate salts react rapidly with aqueous iron(III) chloride to give[Fe(C2O4)3]3−, known asferrioxalate. Othercarboxylate sources, e.g.,citrate andtartrate, bind as well to givecarboxylate complexes. The affinity of iron(III) for oxygen ligands was the basis of qualitative tests for phenols. Although superseded by spectroscopic methods, theferric chloride test is a traditionalcolorimetric test.[26] The affinity of iron(III) for phenols is exploited in theTrinder spot test.[27]
Aqueous iron(III) chloride serves as a one-electron oxidant illustrated by its reaction withcopper(I) chloride to givecopper(II) chloride and iron(II) chloride.
This fundamental reaction is relevant to the use of ferric chloride solutions in etching copper.
The interaction of anhydrous iron(III) chloride withorganolithium andorganomagnesium compounds has been examined often. These studies are enabled because of the solubility of FeCl3 in ethereal solvents, which avoids the possibility of hydrolysis of thenucleophilicalkylating agents. Such studies may be relevant to the mechanism of FeCl3-catalyzedcross-coupling reactions.[28] The isolation of organoiron(III) intermediates requires low-temperature reactions, lest the [FeR4]− intermediates degrade. Usingmethylmagnesium bromide as the alkylation agent, salts of Fe(CH3)4]− have been isolated.[29] Illustrating the sensitivity of these reactions,methyl lithiumLiCH3 reacts with iron(III) chloride to give lithiumtetrachloroferrate(II)Li2[FeCl4]:[30]
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 inferrocene:[31]
This conversion, although not of practical value, was important in the history oforganometallic chemistry where ferrocene is emblematic of the field.[32]
The largest applications of iron(III) chloride aresewage treatment anddrinking water production. By forming highly dispersed networks of Fe-O-Fe containing materials, ferric chlorides serve as coagulant and flocculants.[33] In this application, an aqueous solution ofFeCl3 is treated with base to form afloc ofiron(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 solublephosphate from wastewater.Iron(III) phosphate isinsoluble and thus precipitates as a solid.[34] One potential advantage of its use in water treatment, is that the ferric ion oxidizes (deodorizes)hydrogen sulfide.[35]
It is also used as aleaching agent in chloride hydrometallurgy,[36] for example in the production of Si from FeSi (Silgrain process byElkem).[37]
In another commercial application, a solution of iron(III) chloride is useful for etchingcopper according to the following equation:
The solublecopper(II) chloride is rinsed away, leaving a copper pattern. This chemistry is used in the production ofprinted circuit boards (PCB).[19]
Iron(III) chloride is used in many other hobbies involving metallic objects.[38][39][40][41][42]
In industry, iron(III) chloride is used as a catalyst for the reaction ofethylene withchlorine, forming ethylene dichloride (1,2-dichloroethane):[43]
Ethylene dichloride is acommodity chemical, which is mainly used for the industrial production ofvinyl chloride, themonomer for makingPVC.[44]
Illustrating it use as aLewis acid, iron(III) chloridecatalyseselectrophilic aromatic substitution andchlorinations. In this role, its function is similar to that ofaluminium chloride. In some cases, mixtures of the two are used.[45]
Although iron(III) chlorides are seldom used in practicalorganic synthesis, they have received considerable attention asreagents 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 topolythiophenes upon treatment with ferric chloride.[47] Iron(III) chloride has been shown to promote C-Ccoupling reaction.[48]
Several reagents have been developed based onsupported iron(III) chloride. Onsilica gel, the anhydrous salt has been applied to certaindehydration andpinacol-type rearrangement reactions. A similar reagent but moistened induces hydrolysis orepimerization reactions.[49] Onalumina, ferric chloride has been shown to accelerateene reactions.[50]
When pretreated withsodium hydride, iron(III) chloride gives a hydridereducing agent that convertalkenes andketones intoalkanes andalcohols, respectively.[51]
Iron(III) chloride is a component of useful stains, such asCarnoy's solution, ahistological fixative with many applications. Also, it is used to prepareVerhoeff's stain.[52]
Like many metal halides,FeCl3 naturally occurs as a trace mineral. The rare mineralmolysite is usually associated withvolcanoes andfumaroles.[53][54]
FeCl3-based aerosol are produced by a reaction between iron-rich dust andhydrochloric acid from sea salt. This iron salt aerosol causes about 1–5% of naturally occurring oxidization ofmethane and is thought to have a range of cooling effects; thus, it has been proposed as a catalyst forAtmospheric Methane Removal.[55]
The clouds ofVenus are hypothesized to contain approximately 1%FeCl3 dissolved insulfuric acid.[56][57]
Iron(III) chlorides are widely used in thetreatment of drinking water,[10] so they pose few problems as poisons, at low concentrations.[improper synthesis?] Nonetheless, anhydrous iron(III) chloride, as well as concentratedFeCl3 aqueous solution, is highlycorrosive, and must be handled using proper protective equipment.[20]
_chloride_anhydrate.jpg)
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