Zinc chloride

Zinc chloride
General
Systematic name	Zinc chloride
Other names	Zinc(II) chloride, zinc dichloride, butter of zinc
Molecular formula	ZnCl₂
Molar mass	136.29 g/mol
Appearance	White crystalline solid.
CAS number	^*
Properties
Density and phase	2.907 g/cm³, solid
Solubility in water	432 g/100 mL (25 °C)
in ethanol	100 g/100 mL (12.5 °C)
in acetone, diethyl ether	Soluble
Melting point	275 °C (548 K)
Boiling point	756 °C (1029 K)
Structure
Coordination geometry	Tetrahedral, 4-coordinate, linear in the gas phase.
Crystal structure	Four forms known Hexagonal close-packed (δ) is the only stable form when anhydrous.
Dipole moment	? D
Hazards
MSDS	External MSDS
EU classification	Irritant (I), Corrosive(C).
NFPA 704
R-phrases	, ,
S-phrases	, , , ,
RTECS number	ZH1400000
Supplementary data page
Structure and properties	n, ε_r, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Related compounds
Other anions	Zinc fluoride, zinc bromide, zinc iodide
Other cations	Copper(II) chloride, cadmium chloride
Except where noted otherwise, data are given for materials in their standard state (at 25°C, 100 kPa) Chemical infobox

Zinc chloride is the name of chemical compound Zn Cl₂ or its hydrates. All zinc chlorides are colorless or white and highly soluble in water. ZnCl₂ itself is hygroscopic and can be considered deliquescent. Samples should therefore be protected from sources of moisture, such as the atmosphere.

Four crystalline forms of ZnCl₂ and its hydrates are known, although the anhydrous form appears to exist only in the hexagonal close-packed phase. Rapid cooling of molten ZnCl₂ forms a glassy material.

Concentrated aqueous solutions of zinc chloride have the interesting property of dissolving starch, silk, and cellulose. Thus, solutions of zinc chlorides cannot be filtered through standard filter papers.

Zinc chloride finds wide application in textile processing, metallurgical fluxes and chemical synthesis.

Chemical properties

ZnCl₂ is an ionic solid, although some covalent character is indicated by its low melting point (275 °C). Further evidence for covalency is provided by its high solubility in solvents such as diethyl ether. ZnCl₂ is a mild Lewis acid. Consistent with this character, aqueous solutions of ZnCl₂ have a pH around 4. It is hydrolyzed to an oxychloride when hydrated forms are heated.

In aqueous solution, zinc chloride is a useful source of Zn²⁺ for the preparation of other zinc salts, for example zinc carbonate:

ZnCl₂(aq) + Na₂CO₃(aq) → ZnCO₃(s) + 2 NaCl(aq)

Preparation and purification

Anhydrous ZnCl₂ can be prepared from zinc and hydrogen chloride.

Zn + 2 HCl → ZnCl₂ + H₂

Hydrated forms and aqueous solutions may be readily prepared using concentrated hydrochloric acid and pieces of Zn. Zinc oxide and zinc sulfide react with HCl, without forming hydrogen:

ZnS(s) + 2 HCl(aq) → ZnCl₂(aq) + H₂S(l)

Commercial samples of zinc chloride typically contain water and zinc oxychloride, the main hydrolysis product. Such samples may be purified as follows: 100 g of crude ZnCl₂ are heated to reflux in 800 mL anhydrous dioxan in the presence of zinc metal dust. The mixture is filtered while hot (to remove Zn),and then the filtrate is allowed to cool to give pure ZnCl₂ as a white precipitate. Anhydrous samples can be purified by sublimation in a stream of hydrogen chloride gas, followed by heating to 400 °C in a stream of dry nitrogen gas.

Uses

ZnCl₂ is used as a flux for soldering because of its ability (when molten) to dissolve metal oxides. This property also leads to its use in the manufacture of magnesia cements for dental fillings. ZnCl₂ has also been used as a fireproofing agent, for etching metals, and is also a primary ingredient in fabric refresheners such as Febreze.

In the laboratory, zinc chloride finds wide use, principally as a moderate-strength Lewis acid. It can catalyse (A) the Fischer indole synthesis^{and also (B) Friedel-Crafts acylation reactions involving activated aromatic rings^[10.}

Related to the latter is the classical preparation^{^*} of the dye fluorescein from phthalic anhydride and resorcinol, which involves a Friedel-Crafts acylation. This has in fact been done successfully using the wet ZnCl₂ sample shown in the picture above.

Hydrochloric acid alone reacts poorly with primary alcohols and secondary alcohols, but a combination of HCl with ZnCl₂ (known together as the "Lucas reagent") at 130°C is effective for the preparation of alkyl chlorides. This probably reacts via an S_N2 mechanism with primary alcohols but via S_N1 with secondary alcohols.

Zinc chloride is also able to activate benzylic and allylic halides towards substitution by weak nucleophiles such as alkenes^{^*}:

In similar fashion, ZnCl₂ promotes selective NaBH₃CN reduction of tertiary, allylic or benzylic halides to the corresponding hydrocarbons.

Zinc chloride is also a useful starting point for the synthesis of many organozinc reagents, such as those used in the palladium catalysed Negishi coupling with aryl halides or vinyl halides^{^*}. In such cases the organozinc compound is usually prepared by transmetallation from an organolithium or a Grignard reagent, for example:

Zinc enolates, prepared from alkali metal enolates and ZnCl₂, provide control of stereochemistry in aldol condensation reactions due to chelation on to the zinc. In the example shown below, the threo product was favored over the erythro by a factor of 5:1 when ZnCl₂ in DME/ether was used^{^*}. This is because the chelate is more stable when the bulky phenyl group is pseudo-equatorial rather than pseudo-axial, i.e., threo rather than erythro.

Precautions

Corrosive, irritant. Wear gloves and goggles.

References

N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, 2nd ed., Butterworth-Heinemann, Oxford, UK, 1997.
Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
D. Nicholls, Complexes and First-Row Transition Elements, Macmillan Press, London, 1973.
A. F. Wells, 'Structural Inorganic Chemistry, 5th ed., Oxford University Press, Oxford, UK, 1984.
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.
B. S. Furnell et al., Vogel's Textbook of Practical Organic Chemistry, 5th edition, Longman/Wiley, New York, 1989.
R. L. Shriner, W. C. Ashley, E. Welch, in Organic Syntheses Collective Volume 3, p 725, Wiley, New York, 1955.
(a) S. R. Cooper, in Organic Syntheses Collective Volume 3, p 761, Wiley, New York, 1955. (b) S. Y. Dike, J. R. Merchant, N. Y. Sapre, Tetrahedron, 47, 4775 (1991).
E. Bauml, K. Tschemschlok, R. Pock, H. Mayr, Tetrahedron Letters, 29, 6925 (1988).
S. Kim, Y. J. Kim, K. H. Ahn, Tetrahedron Letters, 24, 3369 (1983).
H. O. House, D. S. Crumrine, A. Y. Teranishi, H. D. Olmstead, Journal of the American Chemical Society, 95, 3310 (1973).

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