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Iron(III) chloride describes the inorganic compounds with the formula Template:Chem2(H2O)x. Also called ferric 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 both hygroscopic. They feature iron in its +3 oxidation state. The anhydrous derivative is a 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 propertiesEdit
All forms of ferric chloride are paramagnetic, 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 per d-orbital. The high spin d5 electronic configuration requires that d-d electronic transitions are spin forbidden, in addition to violating the Laporte rule. This double forbidden-ness results in its solutions being only pale colored. Or, stated more technically, the optical transitions are non-intense. Aqueous ferric sulfate and ferric nitrate, which contain Template:Chem2, are nearly colorless, whereas the chloride solutions are yellow. Thus, the chloride ligands significantly influence the optical properties of the iron center.<ref>Template:Cite book</ref><ref name=Cotton>Template:Cite journal</ref>
StructureEdit
Iron(III) chloride can exist as an anhydrous material and a series of hydrates, which results in distinct structures.
AnhydrousEdit
The anhydrous 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 by transmitted light, they appear purple-red. Anhydrous iron(III) chloride has the [[bismuth(III) iodide|Template:Chem2]] structure, with octahedral Fe(III) centres interconnected by two-coordinate chloride ligands.<ref name="str"/><ref name="UllmannFe" />
Iron(III) chloride has a relatively low melting point and boils at around 315 °C. The vapor consists of the dimer Template:Chem2, much like aluminium chloride. This dimer dissociates into the monomeric Template:Chem2 (with D3h point group molecular symmetry) at higher temperatures, in competition with its reversible decomposition to give iron(II) chloride and chlorine gas.<ref>Template:Cite book</ref>
HydratesEdit
Ferric chloride form hydrates upon exposure to water, reflecting its Lewis acidity. All hydrates exhibit deliquescence, meaning that they become liquid by absorbing moisture from the air. Hydration invariably gives derivatives of aquo complexes with the formula Template:Chem2. This cation can adopt either 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 Template:Chem2, the disesquihydrate Template:Chem2, the trisesquihydrate Template:Chem2, and finally the hexahydrate Template:Chem2. 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.<ref name=Cotton/> The structural formulas are Template:Chem2, Template:Chem2, Template:Chem2, and Template:Chem2. The first three members of this series have the tetrahedral tetrachloroferrate (Template:Chem2) anion.<ref>Template:Cite journal</ref>
SolutionEdit
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Like the solid hydrates, aqueous solutions of ferric chloride also consist of the octahedral Template:Chem2 of unspecified stereochemistry.<ref name=Cotton/> 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 for aquo complexes of a tricationic metal. Iron aquo complexes are prone to olation, the formation of polymeric oxo derivatives. Dilute solutions of ferric chloride produce soluble nanoparticles with molecular weight of 104, which exhibit the property of "aging", i.e., the structure change or evolve over the course of days.<ref>Template:Cite journal</ref> 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.<ref name=ZaaC/>
PreparationEdit
Several hundred tons of anhydrous iron(III) chloride are produced annually. The principal method, called direct chlorination, uses scrap iron as a precursor:<ref name="UllmannFe" />
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.<ref name=UllmannFe/> A similar but laboratory-scale process also has been described.<ref>Template:Cite book</ref><ref name=Brauer>Template:Cite book</ref>
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 by hydrochloric acid followed by chlorination:<ref name="UllmannFe" />
A number of variables apply to these processes, including the oxidation of iron by ferric chloride and the hydration of intermediates.<ref name="UllmannFe" /> 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<ref>Template:Cite book</ref> or trimethylsilyl chloride:<ref>Template:Cite book</ref>
ReactionsEdit
Being high spin d5 electronic configuration iron(III) chlorides are labile, meaning that its Cl- and H2O ligands exchange rapidly with free chloride and water.<ref name="Cotton" /><ref name="greenwood" /> 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 FeCl3Edit
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.<ref name="EROS" />
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:<ref> Template:Cite book</ref>
Alkali metal alkoxides react to give the iron(III) alkoxide complexes. These products have more complicated structures than anhydrous iron(III) chloride.<ref>Template:Cite book</ref><ref>Template:Cite book</ref> In the solid phase a variety of multinuclear complexes have been described for the nominal stoichiometric reaction between Template:Chem2 and sodium ethoxide:
Iron(III) chloride forms a 1:2 adduct with Lewis bases such as triphenylphosphine oxide; e.g., Template:Chem2. The related 1:2 complex Template:Chem2, has been crystallized from ether solution.<ref name=ZaaC>Template:Cite journal</ref>
Iron(III) chloride also reacts with tetraethylammonium chloride to give the yellow salt of the tetrachloroferrate ion (Template:Chem2). Similarly, combining FeCl3 with NaCl and KCl gives Template:Chem2 and Template:Chem2, respectively.<ref>Template:Cite journal</ref>
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.<ref name="UllmannFe" />
In terms of its being an oxidant, iron(III) chloride oxidizes iron powder to form iron(II) chloride via a comproportionation reaction:<ref name=UllmannFe/>
A traditional synthesis of anhydrous ferrous chloride is the reduction of FeCl3 with chlorobenzene:<ref>Template:Cite book</ref>
iron(III) chloride releases chlorine gas when heated above 160 °C, generating ferrous chloride:<ref name=Brauer/>
To suppress this reaction, the preparation of iron(III) chloride requires an excess of chlorinating agent, as discussed above.<ref name=Brauer/><ref name="UllmannFe" />
Hydrated FeCl3Edit
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 Template:Chem2, 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 superseded by spectroscopic methods, the ferric chloride test is a traditional colorimetric test.<ref>Template:Cite book</ref> The affinity of iron(III) for phenols is exploited in the Trinder spot test.<ref>Template:Cite journal</ref>
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.
This fundamental reaction is relevant to the use of ferric chloride solutions in etching copper.
Organometallic chemistryEdit
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 ethereal solvents, which avoids the possibility of hydrolysis of the nucleophilic alkylating agents. Such studies may be relevant to the mechanism of FeCl3-catalyzed cross-coupling reactions.<ref name=Byers>Template:Cite journal</ref> 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.<ref>Template:Cite journal and references therein.</ref> Illustrating the sensitivity of these reactions, methyl lithium Template:Chem2 reacts with iron(III) chloride to give lithium tetrachloroferrate(II) Template:Chem2:<ref>Template:Cite journal</ref>
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.<ref name=EROS/> These diketonate complexes have the advantages that they do not form hydrates, unlike iron(III) chloride, and they are more soluble in relevant solvents.<ref name=Byers/> Cyclopentadienyl magnesium bromide undergoes a complex reaction with iron(III) chloride, resulting in ferrocene:<ref>Template:Cite journal</ref>
This conversion, although not of practical value, was important in the history of organometallic chemistry where ferrocene is emblematic of the field.<ref name = Pauson2001>Template:Cite journal</ref>
UsesEdit
Water treatmentEdit
The largest applications of iron(III) chloride are sewage treatment and drinking water production. By forming highly dispersed networks of Fe-O-Fe containing materials, ferric chlorides serve as coagulant and flocculants.<ref name="wtcbrochure">Template:Cite book</ref> In this application, an aqueous solution of Template:Chem2 is treated with base to form a floc of iron(III) hydroxide (Template:Chem2), also formulated as FeO(OH) (ferrihydrite). This floc facilitates the separation of suspended materials, clarifying the water.<ref name=UllmannFe>Template:Cite book</ref>
Iron(III) chloride is also used to remove soluble phosphate from wastewater. Iron(III) phosphate is insoluble and thus precipitates as a solid.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> One potential advantage of its use in water treatment, is that the ferric ion oxidizes (deodorizes) hydrogen sulfide.<ref>Template:Cite journal</ref>
Etching and metal cleaningEdit
It is also used as a leaching agent in chloride hydrometallurgy,<ref>Template:Cite journal</ref> for example in the production of Si from FeSi (Silgrain process by Elkem).<ref>Template:Cite journal</ref>
In another commercial application, a solution of iron(III) chloride is useful for etching copper according to the following equation:
The soluble copper(II) chloride is rinsed away, leaving a copper pattern. This chemistry is used in the production of printed circuit boards (PCB).<ref name="greenwood">Template:Cite book</ref>
Iron(III) chloride is used in many other hobbies involving metallic objects.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref><ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref><ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref><ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref><ref>Template:Cite book</ref>
Organic chemistryEdit
In industry, iron(III) chloride is used as a catalyst for the reaction of ethylene with chlorine, forming ethylene dichloride (1,2-dichloroethane):<ref>Template:Cite book</ref>
Ethylene dichloride is a commodity chemical, which is mainly used for the industrial production of vinyl chloride, the monomer for making PVC.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>
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.<ref>Template:Cite journal</ref>
Organic synthesis researchEdit
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.<ref name=EROS/> For example, iron(III) chloride oxidizes naphthols to naphthoquinones:<ref name=EROS/><ref>Template:Cite journal</ref> 3-Alkylthiophenes are polymerized to polythiophenes upon treatment with ferric chloride.<ref>Template:Cite journal</ref> Iron(III) chloride has been shown to promote C-C coupling reaction.<ref>Template:Cite journal</ref>
Several reagents have been developed based on supported iron(III) chloride. On silica gel, the anhydrous salt has been applied to certain dehydration and pinacol-type rearrangement reactions. A similar reagent but moistened induces hydrolysis or epimerization reactions.<ref>Template:Cite book</ref> On alumina, ferric chloride has been shown to accelerate ene reactions.<ref>Template:Cite book</ref>
When pretreated with sodium hydride, iron(III) chloride gives a hydride reducing agent that convert alkenes and ketones into alkanes and alcohols, respectively.<ref>Template:Cite book</ref>
HistologyEdit
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.<ref name="ucdavis">Template:Cite encyclopedia</ref>
Natural occurrenceEdit
Like many metal halides, Template:Chem2 naturally occurs as a trace mineral. The rare mineral molysite is usually associated with volcanoes and fumaroles.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref><ref name="IMA">{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>
Template:Chem2-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.<ref>Template:Cite journal</ref>
The clouds of Venus are hypothesized to contain approximately 1% Template:Chem2 dissolved in sulfuric acid.<ref name="kras006">Template:Cite journal</ref><ref>Template:Cite journal</ref>
SafetyEdit
Iron(III) chlorides are widely used in the treatment of drinking water,<ref name="UllmannFe" /> so they pose few problems as poisons, at low concentrations.Template:Synthesis inline Nonetheless, anhydrous iron(III) chloride, as well as concentrated Template:Chem2 aqueous solution, is highly corrosive, and must be handled using proper protective equipment.<ref name=EROS>Template:Cite book</ref>