Zinc chloride

Zinc chloride is an inorganic chemical compound with the formula ZnCl₂·nH₂O, with n ranging from 0 to 4.5, forming hydrates. Zinc chloride, anhydrous and its hydrates, are colorless or white crystalline solids, and are highly soluble in water. Five hydrates of zinc chloride are known, as well as four polymorphs of anhydrous zinc chloride.<ref name="a"/>

All forms of zinc chloride are deliquescent. They can usually be produced by the reaction of zinc or its compounds with some form of hydrogen chloride. Anhydrous zinc compound is a Lewis acid, readily forming complexes with a variety of Lewis bases. Zinc chloride finds wide application in textile processing, metallurgical fluxes, chemical synthesis of organic compounds, such as benzaldehyde, and processes to produce other compounds of zinc.<ref name="a" />

History

Zinc chloride has long been known but currently practiced industrial applications all evolved in the latter half of 20th century.<ref name="a"/>

An amorphous cement formed from aqueous zinc chloride and zinc oxide was first investigated in 1855 by Stanislas Sorel. Sorel later went on to investigate the related magnesium oxychloride cement, which bears his name.<ref>Template:Cite book</ref>

Dilute aqueous zinc chloride was used as a disinfectant under the name "Burnett's Disinfecting Fluid".<ref>Template:Cite book</ref> From 1839 Sir William Burnett promoted its use as a disinfectant as well as a wood preservative. The Royal Navy conducted trials into its use as a disinfectant in the late 1840s, including during the cholera epidemic of 1849; and at the same time experiments were conducted into its preservative properties as applicable to the shipbuilding and railway industries. Burnett had some commercial success with his eponymous fluid. Following his death however, its use was largely superseded by that of carbolic acid and other proprietary products.<ref>Template:Cite journal</ref>

Structure and properties

Unlike other metal dichlorides, zinc dichloride adopts several crystalline forms (polymorphs). Four polymorph are known: α, β, γ, and δ. Each features Template:Chem2 centers surrounded in a tetrahedral manner by four chloride ligands.<ref name=Wells>Template:Cite book</ref>

Form	Crystal system	Pearson symbol	Space group	No.	a (nm)	b (nm)	c (nm)	Z	Density (g/cm3)
α	tetragonal	tI12	ITemplate:Overline2d	122	0.5398	0.5398	0.64223	4	3.00
β	tetragonal	tP6	P42/nmc	137	0.3696	0.3696	1.071	2	3.09
γ	monoclinic	mP36	P21/c	14	0.654	1.131	1.23328	12	2.98
δ	orthorhombic	oP12	Pna21	33	0.6125	0.6443	0.7693	4	2.98

Here a, b, and c are lattice constants, Z is the number of structure units per unit cell, and ρ is the density calculated from the structure parameters.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref><ref>Template:Cite journal</ref>

The orthorhombic form (δ) rapidly changes to another polymorph upon exposure to the atmosphere. A possible explanation is that the Template:Chem2 ions originating from the absorbed water facilitate the rearrangement.<ref name=Wells/> Rapid cooling of molten Template:Chem2 gives a glass.<ref>Template:Cite journal</ref>

Molten Template:Chem2 has a high viscosity at its melting point and a comparatively low electrical conductivity, which increases markedly with temperature.<ref name = "prince">Template:Cite book</ref><ref>Template:Cite book</ref> As indicated by a Raman scattering study, the viscosity is explained by the presence of polymers,.<ref>Template:Cite book</ref> Neutron scattering study indicated the presence of tetrahedral Template:Chem2 centers, which requires aggregation of Template:Chem2 monomers as well.<ref>Template:Cite journal</ref>

Hydrates

A variety of hydrated zinc chloride are known: Template:Chem2 with n = 1, 1.33, 2.5, 3, and 4.5.<ref name=Holleman>Template:Cite book</ref> The 1.33-hydrate, previously thought to be the hemitrihydrate, consists of trans-Zn(H₂O)₄Cl₂ centers with the chlorine atoms connected to repeating ZnCl₄ chains. The hemipentahydrate, structurally formulated [Zn(H₂O)₅][ZnCl₄], consists of Zn(H₂O)₅Cl octahedrons where the chlorine atom is part of a [ZnCl₄]^2- tetrahedera. The trihydrate consists of distinct hexaaquozinc(II) cations and tetrachlorozincate anions; formulated [Zn(H₂O)₆][ZnCl₄]. Finally, the heminonahydrate, structurally formulated [Zn(H₂O)₆][ZnCl₄]·3H₂O also consists of distinct hexaaquozinc(II) cations and tetrachlorozincate anions like the trihydrate but has three extra water molecules. These hydrates can be produced by evaporation of aqueous solutions of zinc chloride at different temperatures.<ref name="1.33">Template:Cite journal</ref><ref name="hydrcrys">Template:Cite journal</ref>

Preparation and purification

Historically, zinc chlorides are prepared from the reaction of hydrochloric acid with zinc metal or zinc oxide. Aqueous acids cannot be used to produce anhydrous zinc chloride. According to an early procedure, a suspension of powdered zinc in diethyl ether is treated with hydrogen chloride, followed by drying<ref>Template:Cite journal</ref> The overall method remains useful in industry, but without the solvent:<ref name="a" />

Template:Chem2

Aqueous solutions may be readily prepared similarly by treating Zn metal, zinc carbonate, zinc oxide, and zinc sulfide with hydrochloric acid:<ref>Template:Cite book</ref>

Template:Chem2

Hydrates can be produced by evaporation of an aqueous solution of zinc chloride. The temperature of the evaporation determines the hydrates. For example, evaporation at room temperature produces the 1.33-hydrate.<ref name="1.33" /><ref name="old">Template:Cite journal</ref> Lower evaporation temperatures produce higher hydrates.<ref name="hydrcrys" />

Commercial samples of zinc chloride typically contain water and products from hydrolysis as impurities. Laboratory samples may be purified by recrystallization from hot dioxane. Anhydrous samples can be purified by sublimation in a stream of hydrogen chloride gas, followed by heating the sublimate to 400 °C in a stream of dry nitrogen gas.<ref name="eros">Template:Cite journal</ref> A simple method relies on treating the zinc chloride with thionyl chloride.<ref>Template:Cite book</ref>

Reactions

Chloride complexes

A number of salts containing the tetrachlorozincate anion, Template:Chem2, are known.<ref name = "prince"/> "Caulton's reagent", Template:Chem2, which is used in organic chemistry, is an example of a salt containing Template:Chem2.<ref>Template:Cite book</ref><ref>Template:Cite journal</ref> The compound Template:Chem2 contains tetrahedral Template:Chem2 and [[Chloride|Template:Chem2]] anions,<ref name=Wells/> so, the compound is not caesium pentachlorozincate, but caesium tetrachlorozincate chloride. No compounds containing the Template:Chem2 ion (hexachlorozincate ion) have been characterized.<ref name=Wells/> The compound Template:Chem2 crystallizes from a solution of Template:Chem2 in hydrochloric acid. It contains a polymeric anion Template:Chem2 with balancing monohydrated hydronium ions, Template:Chem2 ions.<ref name=Wells/>

Adducts

File:CSD CIF TOCMON02.png

Crystal structure of ZnCl₂(thf)₂.<ref>Template:Cite journal</ref>

The adduct with thf Template:Chem2 illustrates the tendency of zinc chloride to form 1:2 adducts with weak Lewis bases. Being soluble in ethers and lacking acidic protons, this complex is used in the synthesis of organozinc compounds.<ref>Template:Cite journal</ref> A related 1:2 complex is Template:Chem2 (zinc dichloride di(hydroxylamine)). Known as Crismer's salt, this complexes releases hydroxylamine upon heating.<ref>Template:Cite book</ref> The distinctive ability of aqueous solutions of Template:Chem2 to dissolve cellulose is attributed to the formation of zinc-cellulose complexes, illustrating the stability of its adducts.<ref>Template:Cite journal</ref> Cellulose also dissolves in molten Template:Chem2 hydrate.<ref>Template:Cite journal</ref> Overall, this behavior is consistent with Zn²⁺ as a hard Lewis acid.

When solutions of zinc chloride are treated with ammonia, diverse ammine complexes are produced. In addition to the tetrahedral 1:2 complex Template:Chem2.<ref>Template:Cite journal</ref><ref>Template:Cite book</ref> the complex Template:Chem2 also has been isolated. The latter contains the Template:Chem2 ion,.<ref name=Wells/> The species in aqueous solution have been investigated and show that Template:Chem2 is the main species present with Template:Chem2 also present at lower Template:Chem2:Zn ratio.<ref>Template:Cite journal</ref>

Aqueous solutions of zinc chloride

Zinc chloride dissolves readily in water to give Template:Chem2 species and some free chloride.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref><ref>Template:Cite journal</ref> Aqueous solutions of Template:Chem2 are acidic: a 6 M aqueous solution has a pH of 1.<ref name=Holleman/> The acidity of aqueous Template:Chem2 solutions relative to solutions of other Zn²⁺ salts (say the sulfate) is due to the formation of the tetrahedral chloro aqua complexes such as [ZnCl₃(H₂O)]⁻.<ref>Template:Cite book</ref> Most metal dichlorides form octahedral complexes, with stronger O-H bonds. The combination of hydrochloric acid and Template:Chem2 gives a reagent known as "Lucas reagent". Such reagents were once used as a test for primary alcohols. Similar reactions are the basis of industrial routes from methanol and ethanol respectively to methyl chloride and ethyl chloride.<ref>Template:Cite journal</ref>

In alkali solution, zinc chloride converts to various zinc hydroxychlorides. These include Template:Chem2, Template:Chem2, Template:Chem2, and the insoluble Template:Chem2. The latter is the mineral simonkolleite.<ref>Template:Cite book {{#invoke:citation/CS1|citation |CitationClass=web }}</ref> When zinc chloride hydrates are heated, hydrogen chloride evolves and hydroxychlorides result.<ref>Template:Cite journal</ref>

In aqueous solution Template:Chem2, as well as other halides (bromide, iodide), behave interchangeably for the preparation of other zinc compounds. These salts give precipitates of zinc carbonate when treated with aqueous carbonate sources:<ref name="a" />

Template:Chem2

Ninhydrin reacts with amino acids and amines to form a colored compound "Ruhemann's purple" (RP). Spraying with a zinc chloride solution, which is colorless, forms a 1:1 complex RP:Template:Chem2, which is more readily detected as it fluoresces more intensely than RP.<ref>Template:Cite book</ref>

Redox

Anhydrous zinc chloride melts and even boils without any decomposition up to 900 °C. When zinc metal is dissolved in molten Template:Chem2 at 500–700 °C, a yellow diamagnetic solution is formed consisting of the Template:Chem2, which has zinc in the oxidation state +1. The nature of this dizinc dication has been confirmed by Raman spectroscopy.<ref name="Holleman"/> Although Template:Chem2 is unusual, mercury, a heavy congener of zinc, forms a wide variety of Template:Chem2 salts.

In the presence of oxygen, zinc chloride oxidizes to zinc oxide above 400 °C. Again, this observation indicates the nonoxidation of Zn²⁺.<ref name="decomp">Template:Cite journal</ref>

Zinc hydroxychloride

Concentrated aqueous zinc chloride dissolves zinc oxide to form zinc hydroxychloride, which is obtained as colorless crystals:<ref>Template:Cite book</ref>

Template:Chem2

The same material forms when hydrated zinc chloride is heated.<ref>Template:Cite book</ref>

The ability of zinc chloride to dissolve metal oxides (MO)<ref name="HOWI">Template:Cite book</ref> is relevant to the utility of Template:Chem2 as a flux for soldering. It dissolves passivating oxides, exposing the clean metal surface.<ref name="HOWI"/>

Organic syntheses with zinc chloride

Zinc chloride is an occasional laboratory reagent often as a Lewis acid. A dramatic example is the conversion of methanol into hexamethylbenzene using zinc chloride as the solvent and catalyst:<ref name = HMB>Template:Cite journal</ref>

Template:Chem2

This kind of reactivity has been investigated for the valorization of C1 precursors.<ref>Template:Cite journal</ref>

Examples of zinc chloride as a Lewis acid include the Fischer indole synthesis:<ref>Template:OrgSynth</ref>

File:ZnCl2 aromatics (cropped indole).gif

Related Lewis-acid behavior is illustrated by a traditional preparation of the dye fluorescein from phthalic anhydride and resorcinol, which involves a Friedel-Crafts acylation.<ref>Template:Cite book</ref> This transformation has in fact been accomplished using even the hydrated Template:Chem2 sample shown in the picture above. Many examples describe the use of zinc chloride in Friedel-Crafts acylation reactions.<ref>Template:OrgSynth</ref><ref>Template:Cite journal</ref>

File:Preparation of Fluorescein.svg

Zinc chloride also activates benzylic and allylic halides towards substitution by weak nucleophiles such as alkenes:<ref>Template:Cite journal</ref>

File:ZnCl2 benzylation (cropped).gif

In similar fashion, Template:Chem2 promotes selective [[sodium cyanoborohydride|Template:Chem2]] reduction of tertiary, allylic or benzylic halides to the corresponding hydrocarbons.<ref name="eros" />

Zinc enolates, prepared from alkali metal enolates and Template:Chem2, provide control of stereochemistry in aldol condensation reactions. This control is attributed to chelation at the zinc. In the example shown below, the threo product was favored over the erythro by a factor of 5:1 when Template:Chem2.<ref>Template:Cite journal</ref>

File:ZnCl2 aldol (cropped).gif

Organozinc precursor

Being inexpensive and anhydrous, ZnCl₂ is a widely used for the synthesis of many organozinc reagents, such as those used in the palladium catalyzed Negishi coupling with aryl halides or vinyl halides. The prominence of this reaction was highlighted by the award of the 2010 Nobel Prize in Chemistry to Ei-ichi Negishi.<ref>Template:Cite journal</ref>

File:ZnCl2 Negishi.gif

Rieke zinc, a highly reactive form of zinc metal, is generated by reduction of zinc dichloride with lithium. Rieke Zn is useful for the preparation of polythiophenes<ref>Template:Cite journal</ref> and for the Reformatsky reaction.<ref>Template:Cite journal</ref>

Uses

Industrial organic chemistry

Zinc chloride is used as a catalyst or reagent in diverse reactions conducted on an industrial scale. Benzaldehyde, 20,000 tons of which is produced annually in Western countries, is produced from inexpensive toluene by exploiting the catalytic properties of zinc dichloride. This process begins with the chlorination of toluene to give benzal chloride. In the presence of a small amount of anhydrous zinc chloride, a mixture of benzal chloride are treated continuously with water according to the following stoichiometry:<ref>Template:Cite book</ref>

Template:Chem2

Similarly zinc chloride is employed in hydrolysis of benzotrichloride, the main route to benzoyl chloride. It serves as a catalyst for the production of methylene-bis(dithiocarbamate).<ref name="a">Template:Ullmann</ref>

As a metallurgical flux

The use of zinc chloride as a flux, sometimes in a mixture with ammonium chloride (see also Zinc ammonium chloride), involves the production of HCl and its subsequent reaction with surface oxides.

Zinc chloride forms two salts with ammonium chloride: Template:Chem2 and Template:Chem2, which decompose on heating liberating HCl, just as zinc chloride hydrate does. The action of zinc chloride/ammonium chloride fluxes, for example, in the hot-dip galvanizing process produces Template:Chem2 gas and ammonia fumes.<ref>Template:Cite book</ref>

Other uses

Relevant to its affinity for these paper and textiles, Template:Chem2 is used as a fireproofing agent and in the process of making Vulcanized fibre, which is made by soaking paper in concentrated zinc chloride.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref> Zinc chloride is also used as a deodorizing agent and to make zinc soaps.<ref name="a" />

Safety and health

Zinc and chloride are essential for life. Zn²⁺ is a component of several enzymes, e.g., carboxypeptidase and carbonic anhydrase. Thus, aqueous solutions of zinc chlorides are rarely problematic as an acute poison.<ref name="a"/> Anhydrous zinc chloride is however an aggressive Lewis acid as it can burn skin and other tissues. Ingestion of zinc chloride, often from soldering flux, requires endoscopic monitoring.<ref>Template:Cite journal</ref> Another source of zinc chloride is zinc chloride smoke mixture ("HC") used in smoke grenades. Containing zinc oxide, hexachloroethane and aluminium powder release zinc chloride, carbon and aluminium oxide smoke, an effective smoke screen.<ref>Template:Cite book</ref> Such smoke screens can lead to fatalities.<ref>Template:Cite book</ref>

Further reading

• N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, 2nd ed., Butterworth-Heinemann, Oxford, UK, 1997.
• Template:RubberBible86th
• The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
• D. Nicholls, Complexes and First-Row Transition Elements, Macmillan Press, London, 1973.
• J. March, Advanced Organic Chemistry, 4th ed., p. 723, Wiley, New York, 1992.
• G. J. McGarvey, in Handbook of Reagents for Organic Synthesis, Volume 1: Reagents, Auxiliaries and Catalysts for C-C Bond Formation, (R. M. Coates, S. E. Denmark, eds.), pp. 220–3, Wiley, New York, 1999.

External links

• Grades and Applications of Zinc Chloride
• PubChem ZnCl₂ summary.