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Zinc chloride

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Zinc chloride
Zinc chloride hydrate
Names
IUPAC name
Zinc chloride
Other names
Zinc(II) chloride
Zinc dichloride
Butter of zinc
Identifiers
ECHA InfoCard 100.028.720 Edit this at Wikidata
EC Number
  • 231-592-0
RTECS number
  • ZH1400000
UN number 2331
Properties
ZnCl2
Molar mass 136.315 g/mol
Appearance white crystalline solid
hygroscopic
Odor odorless
Density 2.907 g/cm3
Melting point 292 °C (558 °F; 565 K)
Boiling point 756 °C (1,393 °F; 1,029 K)
432 g/100 mL (25 °C)
Solubility soluble in ethanol, glycerol and acetone
Solubility in alcohol 430 g/100 mL
Structure
Four forms known: hexagonal δ-form is most stable when anhydrous
Tetrahedral, linear in the gas phase
Hazards
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
3
0
0
Lethal dose or concentration (LD, LC):
350 mg/kg, rat (oral)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Zinc chloride is the name of chemical compound with the formula ZnCl2 and its hydrates. Zinc chlorides, of which nine crystalline forms are known, are colorless or white and highly soluble in water. ZnCl2 itself is hygroscopic and even deliquescent. Samples should therefore be protected from sources of moisture, including the water vapor present in ambient air. Zinc chloride finds wide application in textile processing, metallurgical fluxes, and chemical synthesis. No mineral with this chemical composition is known although a very rare mineral, simonkolleite, Zn5(OH)8Cl2·H2O, is known.

Structure and properties

Four crystalline forms, (polymorphs) , of ZnCl2 are known, and in each case the Zn2+ ions are trigonal planar coordinated to four chloride ions.[1] The pure anhydrous orthorhombic form rapidly changes to one of the other forms on exposure to the atmosphere and a possible explanation is that the presence of OH− facilitates the rearrangement.[1] Rapid cooling of molten ZnCl2 gives a glass, that is, a rigid amorphous solid and this ability has been related to the structure in the melt.[2]

The covalent character of the anhydrous material is indicated by its relatively low melting point of 275 °C.[3] Further evidence for covalency is provided by the high solubility of the dichloride in ethereal solvents where it forms adducts with the formula ZnCl2L2, where L = ligand such as O(C2H5)2. In the gas phase, ZnCl2 molecules are linear with a bond length of 205 pm.[4]

Molten ZnCl2 has a high viscosity at its melting point and a comparatively low electrical conductivity that increases markedly with temperature.[4][5]. A raman spectra study of the melt indicated the presence of polymeric structures [6] and a neutron scattering study indicated the presence of tetrahedral {ZnCl4} complexes.[7]

Hydrates

Five hydrates of zinc chloride are known, ZnCl2(H2O)n where n = 1, 1.5, 2.5, 3 and 4.[8] The tetrahydrate ZnCl2(H2O)4 crystallizes from aqueous solutions of zinc chloride.[8]

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 similarly by treating Zn metal with hydrochloric acid. Zinc oxide and zinc sulfide react with HCl:

ZnS(s) + 2 HCl → ZnCl2(aq) + H2S(g)

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

Commercial samples of zinc chloride typically contain water and products from hydrolysis as impurities. Such 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 submlimate to 400 °C in a stream of dry nitrogen gas. Finally, the simplest method relies on treating the zinc chloride with thionyl chloride.[9]

Reactions

Molten anhydrous ZnCl2 at 500 - 700 °C dissolves zinc metal and on rapid cooling of the melt a yellow diamagnetic glass is formed which Raman studies indicate contain the Zn2+
2 ion.[8]

A number of salts containing the tetrachlorozincate anion, ZnCl2−
4, are known [4]. "Caulton's reagent," V2Cl3(thf)6Zn2Cl6 is an example of a salt containing Zn2Cl2−
6.[10][11] The compound Cs3ZnCl5 contains tetrahedral ZnCl2−
4 and Cl− anions. [1] No compounds containing the ZnCl4−
6 ion have been characterized.[1]

Whilst zinc chloride is very soluble in water, solutions cannot be considered to contain simply solvated Zn2+ ions and Cl− ions, ZnClxH2O(4−x) species are also present.[12][13][14] Aqueous solutions of ZnCl2 are acidic: a 6 M aqueous solution has a pH of 1.[8] The acidity of aqueous ZnCl2 solutions relative to solutions of other Zn2+ salts is due to the formation of the tetrahedral chloro aqua complexes where the reduction in coordination number from 6 to 4 further reduces the strength of the O-H bonds in the solvated water molecules.[15]

In alkali solution in the presence of OH− ion various zinc hydroxychloride anions are present in solution, e.g.ZnOH3Cl2−, ZnOH2Cl2−
2, ZnOHCl2−
3, and Zn5OH2Cl3·H2O (simonkolleite) precipitates.[16]

When ammonia is bubbled through a solution of zinc chloride the hydroxide does not precipitate, instead compounds containing complexed ammonia (ammines) are produced, Zn(NH3)4Cl2·H2O and on concentration ZnCl2(NH3)2.[17] The former contains the Zn(NH3)62+ ion [1] and the latter is molecular with a distorted tetrahedral geometry. [18] The species in aqueous solution have been investigated and show that Zn(NH3)42+ is the main species present with Zn(NH3)3Cl+ also present at lower NH3:Zn ratio. [19]

Aqueous zinc chloride reacts with zinc oxide to form an amorphous cement which was first investigated in the 1855 by Sorel,who later he went on to investigate the related magnesium oxychloride cement which bears his name.[20]

When hydrated zinc chloride is heated, one obtains a residue of Zn(OH)Cl e.g.[21]

ZnCl2·2H2O → ZnCl(OH) + HCl + H2O

The compound ZnCl2·½HCl·H2O may be prepared by careful precipitation from a solution of ZnCl2 acidified with HCl and it contains a polymeric anion (Zn2Cl5 −)n with balancing monohydrated hydronium ions, H5O2+ ions. [1][22]

The formation of highly reactive anhydrous HCl gas formed when zinc chloride hydrates are heated is the basis of qualitative inorganic spot tests.[23]

The use of zinc chloride as a flux, sometimes in a mixture with ammonium chloride, involves the production of HCl and its subsequent reaction with surface oxides. Zinc chloride forms two salts with ammonium chloride, (NH4)ZnCl4 and (NH4)3ClZnCl4, 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 H2 gas and ammonia fumes. [24]

Cellulose dissolves in aqueous solutions of ZnCl2 and zinc-cellulose complexes have been detected.[25]. Cellulose also dissolves in molten ZnCl2 hydrate and carboxylation and acetylation performed on the cellulose polymer.[26]

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

Applications

As a metallurgical flux

Zinc chloride has the ability to attack metal oxides (MO) to give derivatives of the formula MZnOCl2.[citation needed] This reaction is relevant to the utility of ZnCl2 as a flux for soldering — it dissolves oxide coatings exposing the clean metal surface.[8] Fluxes with ZnCl2 as an active ingredient are sometimes called "Tinner's Fluid". Typically this flux was prepared by dissolving zinc foil in dilute hydrochloric acid until the liquid ceased to evolve hydrogen; for this reason, such flux was once known as killed spirits. Because of its corrosive nature, this flux is not suitable for situations where any residue cannot be cleaned away, such as electronic work. This property also leads to its use in the manufacture of 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 Lewis acid. It can catalyse (A) the Fischer indole synthesis[27], and also (B) Friedel-Crafts acylation reactions involving activated aromatic rings[28][29]

Related to the latter is the classical preparation of the dye fluorescein from phthalic anhydride and resorcinol, which involves a Friedel-Crafts acylation.[30] This transformation has in fact been accomplished using even the hydrated 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") is effective for the preparation of alkyl chlorides. Typical reactions are conducted at 130 °C. This reaction probably proceeds via an SN2 mechanism with primary alcohols but SN1 pathway with secondary alcohols.

Zinc chloride also activates benzylic and allylic halides towards substitution by weak nucleophiles such as alkenes[31]:

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 Negishi coupling with aryl halides or vinyl halides.[32] 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.[33] 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 starch, silk, and 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

Smoke bombs

Smoke bombs ("HC") contain zinc oxide and hexachloroethane which when ignited react to form zinc chloride smoke. [34]

Fingerprint detection

Ninhydrin reacts with amino acids and amines to form a coloured compound “Ruhemann’s purple” (RP). Spraying with a zinc chloride solution forms a 1:1complex (RP)ZnCl(H2O)2 which is more readily detected as it fluoresces better than Ruhemann’s purple. [35]

Disinfectant

Historically a dilute aqueous solution of zinc chloride was used as a disinfectant under the name "Burnett's Disinfecting Fluid". [36]

Safety considerations

Zinc chloride is a skin and respiratory irritant according to its MSDS[37]. 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 c d e f A.F. Wells (1984). Structural Inorganic Chemistry. Oxford: Clarendon Press. ISBN 0198553706.
  2. ^ Mackenzie, J. D. (1960). "Structure of Glass-Forming Halides. II. Liquid Zinc Chloride". The Journal of Chemical Physics. 33: 366. doi:10.1063/1.1731151.
  3. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  4. ^ a b c Reg. H. Prince (1994). R. Bruce King (ed.). Encyclopedia of Inorganic Chemistry. John Wiley & Sons. ISBN 0471936200.
  5. ^ H. S. Ray (2006). Introduction to Melts: Molten Salts, Slags and Glasses. Allied Publishers. ISBN 8177648756.
  6. ^ V. Danek (2006). Physico-Chemical Analysis of Molten Electrolytes. Elsevier. ISBN 044452116X.
  7. ^ Price, D L (1991). Journal of Physics Condensed Matter. 3: 9835. doi:10.1088/0953-8984/3/49/001. {{cite journal}}: Missing or empty |title= (help)
  8. ^ a b c d e A. F. Holleman, E. Wiberg (2001). Inorganic Chemistry. Academic Press: San Diego. ISBN 012352651. {{cite book}}: Check |isbn= value: length (help)
  9. ^ A. P. Pray (1990). Inorganic Syntheses. Vol. XXVIII. New York: J. Wiley & Sons. pp. 321–2. ISBN 0471526193. Describes the formation of anhydrous LiCl, CuCl2, ZnCl2, CdCl2, ThCl4, CrCl3, FeCl3, CoCl2, and NiCl2 from the corresponding hydrates.
  10. ^ J. Mulzer, H. Waldmann (1998). Organic Synthesis Highlights III. Wiley-VCH. ISBN 3527295003.
  11. ^ Bouma, Reinder J. (1984). "Identification of the zinc reduction product of VCl3.3THF as [V2Cl3(THF)6]2[Zn2Cl6]". Inorganic Chemistry. 23: 2715. doi:10.1021/ic00185a033.
  12. ^ Irish, D. E. (1963). "Raman Study of Zinc Chloride Solutions". The Journal of Chemical Physics. 39: 3436. doi:10.1063/1.1734212.
  13. ^ Yamaguchi, Toshio (1989). "X-ray diffraction and Raman studies of zinc(II) chloride hydrate melts, ZnCl2.rH2O (r = 1.8, 2.5, 3.0, 4.0, and 6.2)". The Journal of Physical Chemistry. 93: 2620. doi:10.1021/j100343a074.
  14. ^ Pye, Cc; Corbeil, Cr; Rudolph, Ww (2006). "An ab initio investigation of zinc chloro complexes". Physical chemistry chemical physics : PCCP. 8 (46): 5428–36. doi:10.1039/b610084h. ISSN 1463-9076. PMID 17119651. {{cite journal}}: Unknown parameter |month= ignored (help)CS1 maint: multiple names: authors list (link)
  15. ^ I. D. Brown (2006). The Chemical Bond in Inorganic Chemistry: The Bond Valence Model. Oxford University Press. ISBN 0199298815.
  16. ^ X. G. Zhang (1996). Corrosion and Electrochemistry of Zinc. Springer. ISBN 0306453347.
  17. ^ H. T. Vulte (2007). Laboratory Manual of Inorganic Preparations. READ BOOKS. ISBN 1408608405.
  18. ^ Yamaguchi, Toshio (1981). "The Crystal Structure of Diamminedichlorozinc(II), ZnCl2(NH3)2. A New Refinement". Acta Chemica Scandinavica. 35a: 727. doi:10.3891/acta.chem.scand.35a-0727.
  19. ^ Yamaguchi, Toshio (1978). "X-Ray Diffraction Studies on the Structures of Tetraammine- and Triamminemonochlorozinc(II) Ions in Aqueous Solution". Bulletin of the Chemical Society of Japan. 51: 3227. doi:10.1246/bcsj.51.3227.
  20. ^ Alan D. Wilson, John W. Nicholson (1993). Acid-base Cements: Their Biomedical and Industrial Applications. Cambridge University Press. ISBN 0521372224.
  21. ^ J. E. House (2008). Inorganic Chemistry. Academic Press. ISBN 0123567866.
  22. ^ Joseph William Mellor (1946). A Comprehensive Treatise on Inorganic and Theoretical Chemistry. Longmans, Green.
  23. ^ Feigl, Fritz (1956). "Some applications of fusion reactions with zinc chloride in inorganic spot test analysis". Microchimica Acta. 44: 1310. doi:10.1007/BF01257465.
  24. ^ American Society for Metals (1990). ASM handbook. ASM International. ISBN 0871700212.
  25. ^ Qin Xu , Li-Fu Chen (1999). "Ultraviolet spectra and structure of zinc-cellulose complexes in zinc chloride solution". Journal of Applied Polymer Science. 71: 1441–1446. doi:10.1002/(SICI)1097-4628(19990228)71:9<1441::AID-APP8>3.0.CO;2-G.
  26. ^ Fischer, S. (2003). Cellulose. 10: 227. doi:10.1023/A:1025128028462. {{cite journal}}: Missing or empty |title= (help)
  27. ^ R. L. Shriner, W. C. Ashley, E. Welch (1955). Organic Syntheses Collective Volume 3. Wiley, New York. p. 725.{{cite book}}: CS1 maint: multiple names: authors list (link)
  28. ^ S. R. Cooper (1955). Organic Syntheses Collective. Vol. 3. Wiley, New York. p. 761.
  29. ^ S. Y. Dike, J. R. Merchant, N. Y. Sapre (1991). "A new and efficient general method for the synthesis of 2-spirobenzopyrans: First synthesis of cyclic analogues of precocene I and related compounds". Tetrahedron. 47: 4775. doi:10.1016/S0040-4020(01)86481-4.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  30. ^ B. S. Furnell; et al. (1989). Vogel's Textbook of Practical Organic Chemistry, 5th edition. Longman/Wiley, New York. {{cite book}}: Explicit use of et al. in: |author= (help)
  31. ^ E. Bauml, K. Tschemschlok, R. Pock, H. Mayr (1988). "Synthesis of γ-lactones from alkenes employing p-methoxybenzyl chloride as +CH2_CO−2 equivalent". Tetrahedron Letters. 29: 6925. doi:10.1016/S0040-4039(00)88476-2.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  32. ^ S. Kim, Y. J. Kim, K. H. Ahn (1983). Tetrahedron Letters. 24: 3369. {{cite journal}}: Missing or empty |title= (help)CS1 maint: multiple names: authors list (link)
  33. ^ H. O. House, D. S. Crumrine, A. Y. Teranishi, H. D. Olmstead (1973). "Chemistry of carbanions. XXIII. Use of metal complexes to control the aldol condensation". Journal of the American Chemical Society. 95: 3310. doi:10.1021/ja00791a039.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  34. ^ B. E. Sample (1997). Methods for Field Studies of Effects of Military Smokes, Obscurants, and Riot-control Agents on Threatened and Endangered Species. DIANE Publishing. ISBN 1428912339.
  35. ^ E. Roland Menzel (1999). Fingerprint Detection with Lasers. CRC Press. ISBN 0824719743.
  36. ^ H. Watts (1869). A Dictionary of Chemistry and the Allied Branches of Other Sciences,. Longmans, Green.
  37. ^ http://www.jtbaker.com/msds/englishhtml/z2280.htm

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. J. March, Advanced Organic Chemistry, 4th ed., p. 723, Wiley, New York, 1992.
  6. 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.
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