Zinc chloride

Zinc chloride is the name of metallurgical fluxes, and chemical synthesis.

Structure and basic properties

Four crystalline forms, so-called hydrates and at least one mixed hydroxide, ZnClOH.^[1]

The pH of 1.^[2]

Four hydrates of zinc chloride are known. ZnCl₂(H₂O)₄ crystallizes from aqueous solutions of zinc chloride. Also characterized are ZnCl₂(H₂O)_n where n = 1, 1.5, 2.5, and 3.^[2] When hydrated zinc chloride is heated, one obtains a residue of ZnOHCl.

In aqueous solution, zinc chloride fully dissociated into Zn²⁺. Thus, although many zinc salts have different formulas and different crystal structures, these salts behave very similarly in aqueous solution. For example, solutions prepared from any of the polymorphs of ZnCl₂ as well as other halides (bromide, iodide) and the sulfate can often be used interchangeably for the preparation of other zinc compounds. Illustrative is the preparation of zinc carbonate:

ZnCl₂(aq) + NaCl(aq)

Preparation and purification

Anhydrous ZnCl₂ can be prepared from hydrogen chloride.

Zn + 2 HCl → ZnCl₂ + H₂

Hydrated forms and aqueous solutions may be readily prepared by treating pieces of Zn metal with concentrated hydrochloric acid. Zinc oxide and zinc sulfide react with HCl:

ZnS(g)

Unlike many other elements, zinc essentially exists in only one oxidation state, 2+, which simplifies purification.

Commercial samples of zinc chloride typically contain water and products from thionyl chloride.^[3]

Applications

As a metallurgical flux

Zinc chloride has the ability to attack metal oxides (MO) to give derivatives of the formula MZnOCl₂. This reaction is relevant to the utility of ZnCl₂ as a magnesia cements for dental fillings and certain mouthwashes as an active ingredient.

In organic synthesis

In the laboratory, zinc chloride finds wide use, principally as a moderate-strength Friedel-Crafts acylation reactions involving activated aromatic rings^[5][6]

Related to the latter is the classical preparation of the dye fluorescein from Friedel-Crafts acylation.^[7] This transformation has in fact been accomplished using even the hydrated ZnCl₂ sample shown in the picture above.

S_N1 pathway with secondary alcohols.

Zinc chloride also activates benzylic and allylic halides towards substitution by weak alkenes^[8]:

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 reagent for the synthesis of many organozinc reagents, such as those used in the palladium catalysed Grignard reagent, for example:

Zinc ether was used.^[10] The chelate is more stable when the bulky phenyl group is pseudo-equatorial rather than pseudo-axial, i.e., threo rather than erythro.

In textile processing

Concentrated aqueous solutions of zinc chloride (more than 64% weight/weight zinc chloride in water) have the interesting property of dissolving cellulose. Thus, such solutions cannot be filtered through standard filter papers. Relevant to its affinity for these materials, ZnCl₂ is used as a fireproofing agent and in fabric "refresheners" such as Febreze

Safety considerations

Zinc salts are relatively non-toxic. Precautions that apply to anhydrous ZnCl₂ are those applicable to other anhydrous metal halides, i.e. hydrolysis can be exothermic and contact should be avoided. Concentrated solutions are acidic and corrosive and specifically attack cellulose and silk as Lewis acids. See MSDS in table.

References

1. ^ ^{a b} Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
2. ^ ^{a b c} Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
3. ^ Pray, A. P. “Anhydrous Metal Chlorides” "Inorganic Syntheses," vol. XXVIII, 321-2, 1990ISBN 0-471-52619-3. Describes the formation of anhydrous LiCl, CuCl₂, ZnCl₂, CdCl₂, ThCl₄, CrCl₃, FeCl₃, CoCl₂, and NiCl₂ from the corresponding hydrates.
4. ^ R. L. Shriner, W. C. Ashley, E. Welch, in Organic Syntheses Collective Volume 3, p 725, Wiley, New York, 1955.
5. ^ S. R. Cooper, in Organic Syntheses Collective Volume 3, p 761, Wiley, New York, 1955.
6. ^ S. Y. Dike, J. R. Merchant, N. Y. Sapre, Tetrahedron, 47, 4775 (1991)
7. ^ B. S. Furnell et al., Vogel's Textbook of Practical Organic Chemistry, 5th edition, Longman/Wiley, New York, 1989.
8. ^ E. Bauml, K. Tschemschlok, R. Pock, H. Mayr, Tetrahedron Letters, 29, 6925 (1988)
9. ^ S. Kim, Y. J. Kim, K. H. Ahn, Tetrahedron Letters, 24, 3369 (1983).
10. ^ H. O. House, D. S. Crumrine, A. Y. Teranishi, H. D. Olmstead, Journal of the American Chemical Society, 95, 3310 (1973)

Bibliography

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.