Zinc chloride

Overview

Zinc chloride (ZnCl₂) is a colorless or white compound of zinc and chlorine that is extremely hygroscopic.

Four crystalline structures have been reported, although in pure form (i.e. water-free) only the delta (hexagonal close-packed) phase can form. It can be quenched from the melt to form a glassy material.

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

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

Chemical properties

Zinc chloride is an ionic salt, though some covalent character is indicated by its low melting point (275 °C) and its high solubility in solvents such as diethyl ether. It behaves as a mild Lewis acid, and aqueous solutions 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 & purification

Anhydrous zinc chloride can be prepared from zinc and hydrogen chloride.

Zn + 2 HCl(g) → ZnCl₂(s) + H₂(g)

Hydrated forms and aqueous solutions may be readily prepared using standard acid-base methods, or from one of its ores, zinc sulfide:

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), then 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 up to 400 °C in a stream of dry nitrogen gas.

Uses

One use for zinc chloride is as a flux for soldering. This is 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 and for etching metals.

In the laboratory, zinc chloride finds wide use, principally as a moderate-strength Lewis acid. It can catalyse (A) the Fischer indole synthesis^[9], and also (B) Friedel-Crafts acylation reactions involving activated aromatic rings ^[10].

Related to the latter is the classical preparation^[8] 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 alcohol]s 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^[11]:

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 ^[12]. 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^[13]. 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.

Suppliers/Manufacturers

• Alfa
• Sigma-Aldrich
• Strem
• Fisher
• VWR

Template: inorganic_stylesheet1

References

1. N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, 2nd ed., Butterworth-Heinemann, Oxford, UK, 1997.
2. Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
3. The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
4. D. Nicholls, Complexes and First-Row Transition Elements, Macmillan Press, London, 1973.
5. A. F. Wells, 'Structural Inorganic Chemistry, 5th ed., Oxford University Press, Oxford, UK, 1984.
6. J. March, Advanced Organic Chemistry, 4th ed., p. 723, Wiley, New York, 1992.
7. 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.
8. B. S. Furnell et al., Vogel's Textbook of Practical Organic Chemistry, 5th edition, Longman/Wiley, New York, 1989.
9. R. L. Shriner, W. C. Ashley, E. Welch, in Organic Syntheses Collective Volume 3, p 725, Wiley, New York, 1955.
10. (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).
11. E. Bauml, K. Tschemschlok, R. Pock, H. Mayr, Tetrahedron Letters, 29, 6925 (1988).
12. S. Kim, Y. J. Kim, K. H. Ahn, Tetrahedron Letters, 24, 3369 (1983).
13. H. O. House, D. S. Crumrine, A. Y. Teranishi, H. D. Olmstead, Journal of the American Chemical Society, 95, 3310 (1973).