Zinc chloride: Difference between revisions
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'''Zinc chloride''' is an [[Inorganic chemistry|inorganic]] [[chemical compound]] with the [[chemical formula|formula]] ZnCl<sub>2</sub>·''n''H<sub>2</sub>O, with ''n'' ranging from 0 to 4.5, forming [[water of hydration|hydrates]]. Zinc chloride, anhydrous and its hydrates, are colorless or white [[crystalline]] solids, and are highly [[Solubility|soluble]] in [[water]]. Five hydrates of zinc chloride are known, as well as four forms of anhydrous zinc chloride.<ref name="a"/> All forms of zinc chloride are [[deliquescent]]. Zinc chloride finds wide application in [[textile]] processing, [[flux (metallurgy)|metallurgical fluxes]], and chemical synthesis.<ref name="a"/> In a major monograph, zinc chlorides have been described as "one of the important compounds of zinc."<ref>{{Greenwood&Earnshaw2nd|page = 1211}}</ref> |
'''Zinc chloride''' is an [[Inorganic chemistry|inorganic]] [[chemical compound]] with the [[chemical formula|formula]] ZnCl<sub>2</sub>·''n''H<sub>2</sub>O, with ''n'' ranging from 0 to 4.5, forming [[water of hydration|hydrates]]. Zinc chloride, anhydrous and its hydrates, are colorless or white [[crystalline]] solids, and are highly [[Solubility|soluble]] in [[water]]. Five hydrates of zinc chloride are known, as well as four forms of anhydrous zinc chloride.<ref name="a"/> |
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All forms of zinc chloride are [[deliquescent]]. Zinc chloride finds wide application in [[textile]] processing, [[flux (metallurgy)|metallurgical fluxes]], and chemical synthesis.<ref name="a" /> In a major [[monograph]], zinc chlorides have been described as "one of the important compounds of zinc."<ref>{{Greenwood&Earnshaw2nd|page = 1211}}</ref> |
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⚫ | 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 [[Sorel cement|magnesium oxychloride cement]], which bears his name.<ref>{{cite book |author1=Wilson, A. D. |author2=Nicholson, J. W. | year = 1993 | title = Acid-Base Cements: Their Biomedical and Industrial Applications | publisher = Cambridge University Press | isbn = 978-0-521-37222-0 }}</ref> |
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⚫ | Dilute aqueous zinc chloride was used as a disinfectant under the name "Burnett's Disinfecting Fluid".<ref>{{cite book | author = Watts, H. | year = 1869 | title = A Dictionary of Chemistry and the Allied Branches of Other Sciences | publisher = Longmans, Green | url = https://archive.org/details/adictionarychem11wattgoog }}</ref> From 1839 [[Sir William Burnett]] promoted its use as a disinfectant as well as a wood preservative.<ref>{{cite journal |last1=McLean |first1=David |title=Protecting wood and killing germs: 'Burnett's Liquid' and the origins of the preservative and disinfectant industries in early Victorian Britain |journal=Business History |date=April 2010 |volume=52 |issue=2 |pages=285–305|doi=10.1080/00076791003610691 |s2cid=154790730 }}</ref> The Royal Navy conducted trials into its use as a disinfectant in the late 1840s, including during the [[1846–1860 cholera pandemic|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.{{Citation needed|date=June 2024}} |
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==Structure and properties== |
==Structure and properties== |
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The orthorhombic form (δ) rapidly changes to one of the other forms on exposure to the atmosphere. A possible explanation is that the {{chem2|OH−}} ions originating from the absorbed water facilitate the rearrangement.<ref name=Wells/> Rapid cooling of molten {{chem2|ZnCl2}} gives a [[glass]].<ref>{{cite journal |author1=Mackenzie, J. D. |author2=Murphy, W. K. | title = Structure of Glass-Forming Halides. II. Liquid Zinc Chloride | journal = The Journal of Chemical Physics | year = 1960 | volume = 33 | issue = 2 | pages = 366–369 | doi = 10.1063/1.1731151 |bibcode=1960JChPh..33..366M }}</ref> |
The orthorhombic form (δ) rapidly changes to one of the other forms on exposure to the atmosphere. A possible explanation is that the {{chem2|OH−}} ions originating from the absorbed water facilitate the rearrangement.<ref name=Wells/> Rapid cooling of molten {{chem2|ZnCl2}} gives a [[glass]].<ref>{{cite journal |author1=Mackenzie, J. D. |author2=Murphy, W. K. | title = Structure of Glass-Forming Halides. II. Liquid Zinc Chloride | journal = The Journal of Chemical Physics | year = 1960 | volume = 33 | issue = 2 | pages = 366–369 | doi = 10.1063/1.1731151 |bibcode=1960JChPh..33..366M }}</ref> |
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Molten {{chem2|ZnCl2}} has a high viscosity at its melting point and a comparatively low electrical conductivity, which increases markedly with temperature.<ref name = "prince">{{cite book | author = Prince, R. H. | year = 1994 | title = Encyclopedia of Inorganic Chemistry | editor = King, R. B. | publisher = John Wiley & Sons | isbn = 978-0-471-93620-6 }}</ref><ref>{{cite book | author = Ray, H. S. | year = 2006| title = Introduction to Melts: Molten Salts, Slags and Glasses | publisher = Allied Publishers | isbn = 978-81-7764-875-1 }}</ref> As indicated by a [[Raman scattering]] study, the viscosity is explained by the presence of polymers,<ref>{{cite book | author = Danek, V. | year = 2006 | title = Physico-Chemical Analysis of Molten Electrolytes | publisher = Elsevier | isbn = 978-0-444-52116-3 }}</ref> |
Molten {{chem2|ZnCl2}} has a high viscosity at its melting point and a comparatively low electrical conductivity, which increases markedly with temperature.<ref name = "prince">{{cite book | author = Prince, R. H. | year = 1994 | title = Encyclopedia of Inorganic Chemistry | editor = King, R. B. | publisher = John Wiley & Sons | isbn = 978-0-471-93620-6 }}</ref><ref>{{cite book | author = Ray, H. S. | year = 2006| title = Introduction to Melts: Molten Salts, Slags and Glasses | publisher = Allied Publishers | isbn = 978-81-7764-875-1 }}</ref> As indicated by a [[Raman scattering]] study, the viscosity is explained by the presence of polymers,.<ref>{{cite book | author = Danek, V. | year = 2006 | title = Physico-Chemical Analysis of Molten Electrolytes | publisher = Elsevier | isbn = 978-0-444-52116-3 }}</ref> [[Neutron scattering]] study indicated the presence of tetrahedral {{chem2|ZnCl4}} centers, which requires aggregation of {{chem2|ZnCl2}} monomers as well.<ref>{{cite journal | last1 = Price | first1 = D. L. | last2 = Saboungi | first2 = M.-L. | last3 = Susman | first3 = S. | last4 = Volin | first4 = K. J. | last5 = Wright | first5 = A. C. | title = Neutron Scattering Function of Vitreous and Molten Zinc Chloride | journal = Journal of Physics: Condensed Matter | year = 1991 | volume = 3 | issue = 49 | pages = 9835–9842 | doi = 10.1088/0953-8984/3/49/001 | bibcode = 1991JPCM....3.9835P | s2cid = 250902741 }}</ref> |
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===Hydrates=== |
===Hydrates=== |
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==Preparation and purification== |
==Preparation and purification== |
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Historically, anhydrous zinc chloride |
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>{{cite journal |doi=10.1039/JR9320002282 |title=Notes: The Preparation of Pure Zinc Chloride |date=1932 |last1=Hamilton |first1=R. T. |last2=Butler |first2=J. A. V. |journal=Journal of the Chemical Society (Resumed) |pages=2283–4 }}</ref> The overall method remains useful in industry, but without the solvent:<ref name="a" /> |
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: {{chem2|Zn + 2 HCl → ZnCl2 + H2}} |
: {{chem2|Zn + 2 HCl → ZnCl2 + H2}} |
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==Reactions== |
==Reactions== |
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===Chloride complexes=== |
===Chloride complexes=== |
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A number of salts containing the [[tetrachlorozincate]] anion, {{chem2|[ZnCl4](2−)}}, are known.<ref name = "prince"/> "Caulton's reagent", {{chem2|[[Vanadium|V]]2Cl3([[thf]])6] [Zn2Cl6]}}, which is used in organic chemistry, is an example of a salt containing {{chem2|[Zn2Cl6](2−)}}.<ref>{{cite book | volume = 3 |editor1=Mulzer, J. |editor2=Waldmann, H. | title = Organic Synthesis Highlights | year = 1998 | publisher = Wiley-VCH | isbn = 978-3-527-29500-5 }}</ref><ref>{{cite journal | last1 = Bouma | first1 = R. J. | last2 = Teuben | first2 =J. H. | last3 = Beukema | first3 = W. R. | last4 = Bansemer | first4 = R. L. | last5 = Huffman | first5 = J. C. | last6 = Caulton | first6 = K. G. | title = Identification of the Zinc Reduction Product of VCl<sub>3</sub> · 3THF as <nowiki>[</nowiki>V<sub>2</sub>Cl<sub>3</sub>(THF)<sub>6</sub><nowiki>]</nowiki><sub>2</sub><nowiki>[</nowiki>Zn<sub>2</sub>Cl<sub>6</sub><nowiki>]</nowiki> | journal = Inorganic Chemistry | year = 1984 | volume = 23 | issue = 17 | pages = 2715–2718 | doi = 10.1021/ic00185a033 }}</ref> The compound {{chem2|[[caesium|Cs]]3ZnCl5}} contains [[Tetrahedral molecular geometry|tetrahedral]] {{chem2|[ZnCl4](2−)}} and [[Chloride|{{chem2|Cl−}}]] anions,<ref name=Wells/> so, the compound is not caesium pentachlorozincate, but caesium tetrachlorozincate chloride. No compounds containing the {{chem2|[ZnCl6](4−)}} ion (hexachlorozincate ion) have been characterized.<ref name=Wells/> The compound {{chem2|ZnCl2*0.5HCl*H2O}} |
A number of salts containing the [[tetrachlorozincate]] anion, {{chem2|[ZnCl4](2−)}}, are known.<ref name = "prince"/> "Caulton's reagent", {{chem2|[[Vanadium|V]]2Cl3([[thf]])6] [Zn2Cl6]}}, which is used in organic chemistry, is an example of a salt containing {{chem2|[Zn2Cl6](2−)}}.<ref>{{cite book | volume = 3 |editor1=Mulzer, J. |editor2=Waldmann, H. | title = Organic Synthesis Highlights | year = 1998 | publisher = Wiley-VCH | isbn = 978-3-527-29500-5 }}</ref><ref>{{cite journal | last1 = Bouma | first1 = R. J. | last2 = Teuben | first2 =J. H. | last3 = Beukema | first3 = W. R. | last4 = Bansemer | first4 = R. L. | last5 = Huffman | first5 = J. C. | last6 = Caulton | first6 = K. G. | title = Identification of the Zinc Reduction Product of VCl<sub>3</sub> · 3THF as <nowiki>[</nowiki>V<sub>2</sub>Cl<sub>3</sub>(THF)<sub>6</sub><nowiki>]</nowiki><sub>2</sub><nowiki>[</nowiki>Zn<sub>2</sub>Cl<sub>6</sub><nowiki>]</nowiki> | journal = Inorganic Chemistry | year = 1984 | volume = 23 | issue = 17 | pages = 2715–2718 | doi = 10.1021/ic00185a033 }}</ref> The compound {{chem2|[[caesium|Cs]]3ZnCl5}} contains [[Tetrahedral molecular geometry|tetrahedral]] {{chem2|[ZnCl4](2−)}} and [[Chloride|{{chem2|Cl−}}]] anions,<ref name=Wells/> so, the compound is not caesium pentachlorozincate, but caesium tetrachlorozincate chloride. No compounds containing the {{chem2|[ZnCl6](4−)}} ion (hexachlorozincate ion) have been characterized.<ref name=Wells/> The compound {{chem2|ZnCl2*0.5HCl*H2O}} crystallizes from a solution of {{chem2|ZnCl2}} in [[hydrochloric acid]]. It contains a polymeric anion {{chem2|(Zn2Cl5−)_{''n''}|}} with balancing monohydrated [[hydronium]] ions, {{chem2|H5O2+}} ions.<ref name=Wells/><!--<ref>{{cite book | author-link = Joseph William Mellor | author = Mellow, J. W. | year = 1946 | title = A Comprehensive Treatise on Inorganic and Theoretical Chemistry | publisher = Longmans, Green }}</ref>--> |
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===Adducts=== |
===Adducts=== |
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[[File:CSD CIF TOCMON02.png|thumb|left|Crystal structure of ZnCl<sub>2</sub>(thf)<sub>2</sub>.<ref>{{cite journal |doi=10.1016/j.mcat.2023.113393 |title=Reactivity of niobium(V) pentaalkoxide complexes: Ring-opening metathesis polymerization of norbornene |date=2023 |last1=Nagata |first1=Tatsuki |last2=Aratani |first2=Shunsuke |last3=Nomura |first3=Moegi |last4=Fuji |first4=Maito |last5=Sotani |first5=Taichi |last6=Sogawa |first6=Hiromitsu |last7=Sanda |first7=Fumio |last8=Yajima |first8=Tatsuo |last9=Obora |first9=Yasushi |journal=Molecular Catalysis |volume=547 }}</ref>]] |
[[File:CSD CIF TOCMON02.png|thumb|left|Crystal structure of ZnCl<sub>2</sub>(thf)<sub>2</sub>.<ref>{{cite journal |doi=10.1016/j.mcat.2023.113393 |title=Reactivity of niobium(V) pentaalkoxide complexes: Ring-opening metathesis polymerization of norbornene |date=2023 |last1=Nagata |first1=Tatsuki |last2=Aratani |first2=Shunsuke |last3=Nomura |first3=Moegi |last4=Fuji |first4=Maito |last5=Sotani |first5=Taichi |last6=Sogawa |first6=Hiromitsu |last7=Sanda |first7=Fumio |last8=Yajima |first8=Tatsuo |last9=Obora |first9=Yasushi |journal=Molecular Catalysis |volume=547 }}</ref>]] |
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The adduct with thf {{chem2|ZnCl2(thf)2}} illustrates the tendency of zinc chloride to form 1:2 adducts with weak [[Lewis base]]s. Being soluble in ethers and lacking acidic protons, this complex is used in the synthesis of [[organozinc compound]]s.<ref>{{cite journal |doi=10.1002/zaac.19976230163 |title=Difluorenylzink als Alkylierungsmittel zur Darstellung von Triorganometallanen der 13. Gruppe. Synthese und Kristallstruktur von [GaFl<sub>3</sub>(THF)] · Toluol (Fl = Fluorenyl) |date=1997 |last1=Dashti |first1=Anahita |last2=Niediek |first2=Katharina |last3=Werner |first3=Bert |last4=Neumüller |first4=Bernhard |journal=Zeitschrift für Anorganische und Allgemeine Chemie |volume=623 |issue=1–6 |pages=394–402 }}</ref> A related 1:2 complex is {{chem2|ZnCl2(NH2OH)2}} (zinc dichloride di(hydroxylamine)). Known as Crismer's salt, this complexes releases [[hydroxylamine]] upon heating.<ref>{{cite book |doi=10.1002/9780470132401.ch2|chapter=Dichlorobis(hydroxylamine)zinc(II) (Crismer's Salt)|year=1967|volume=9|last1=Walker|first1=John E.|last2=Howell|first2=David M.|title=Inorganic Syntheses|pages=2–3|isbn=978-0-470-13240-1}}</ref> The |
The adduct with thf {{chem2|ZnCl2(thf)2}} illustrates the tendency of zinc chloride to form 1:2 adducts with weak [[Lewis base]]s. Being soluble in ethers and lacking acidic protons, this complex is used in the synthesis of [[organozinc compound]]s.<ref>{{cite journal |doi=10.1002/zaac.19976230163 |title=Difluorenylzink als Alkylierungsmittel zur Darstellung von Triorganometallanen der 13. Gruppe. Synthese und Kristallstruktur von [GaFl<sub>3</sub>(THF)] · Toluol (Fl = Fluorenyl) |date=1997 |last1=Dashti |first1=Anahita |last2=Niediek |first2=Katharina |last3=Werner |first3=Bert |last4=Neumüller |first4=Bernhard |journal=Zeitschrift für Anorganische und Allgemeine Chemie |volume=623 |issue=1–6 |pages=394–402 }}</ref> A related 1:2 complex is {{chem2|ZnCl2(NH2OH)2}} (zinc dichloride di(hydroxylamine)). Known as Crismer's salt, this complexes releases [[hydroxylamine]] upon heating.<ref>{{cite book |doi=10.1002/9780470132401.ch2|chapter=Dichlorobis(hydroxylamine)zinc(II) (Crismer's Salt)|year=1967|volume=9|last1=Walker|first1=John E.|last2=Howell|first2=David M.|title=Inorganic Syntheses|pages=2–3|isbn=978-0-470-13240-1}}</ref> The distinctive ability of aqueous solutions of {{chem2|ZnCl2}} to dissolve [[cellulose]] is attributed to the formation of zinc-cellulose complexes, illustrating the stability of its adducts.<ref>{{cite journal |author1=Xu, Q. |author2=Chen, L.-F. | title = Ultraviolet Spectra and Structure of Zinc-Cellulose Complexes in Zinc Chloride Solution | journal = Journal of Applied Polymer Science | year = 1999 | volume = 71 | issue = 9 | pages = 1441–1446 | doi = 10.1002/(SICI)1097-4628(19990228)71:9<1441::AID-APP8>3.0.CO;2-G }}</ref> Cellulose also dissolves in molten {{chem2|ZnCl2}} hydrate.<ref>{{cite journal | last1 = Fischer | first1 = S. | last2 = Leipner | first2 = H. | last3 = Thümmler | first3 = K. | last4 = Brendler | first4 = E. | last5 = Peters | first5 = J. | title = Inorganic Molten Salts as Solvents for Cellulose | journal = Cellulose | year = 2003 | volume = 10 | issue = 3 | pages = 227–236 | doi = 10.1023/A:1025128028462 | s2cid = 92194004 }}</ref> Overall, this behavior is consistent with Zn<sup>2+</sup> as a [[HSAB theory|hard]] Lewis acid. |
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When solutions of zinc chloride are treated with [[ammonia]], diverse [[transition metal ammine complex|ammine]] complexes are produced. In addition to the tetrahedral 1:2 complex {chem2|ZnCl2(NH3)2}}.<ref>{{cite journal | last1 = Yamaguchi | first1 = T. | last2 = Lindqvist | first2 = O. | title = The Crystal Structure of Diamminedichlorozinc(II), ZnCl<sub>2</sub>(NH<sub>3</sub>)<sub>2</sub>. A New Refinement | journal = Acta Chemica Scandinavica A | year = 1981 | volume = 35 | issue = 9 | pages = 727–728 | doi = 10.3891/acta.chem.scand.35a-0727 | url = http://actachemscand.org/pdf/acta_vol_35a_p0727-0728.pdf | doi-access = free }}</ref><ref>{{cite book | author = Vulte, H. T. | title = Laboratory Manual of Inorganic Preparations | publisher = Read Books | year = 2007 | isbn = 978-1-4086-0840-1 }}</ref> |
When solutions of zinc chloride are treated with [[ammonia]], diverse [[transition metal ammine complex|ammine]] complexes are produced. In addition to the tetrahedral 1:2 complex {{chem2|ZnCl2(NH3)2}}.<ref>{{cite journal | last1 = Yamaguchi | first1 = T. | last2 = Lindqvist | first2 = O. | title = The Crystal Structure of Diamminedichlorozinc(II), ZnCl<sub>2</sub>(NH<sub>3</sub>)<sub>2</sub>. A New Refinement | journal = Acta Chemica Scandinavica A | year = 1981 | volume = 35 | issue = 9 | pages = 727–728 | doi = 10.3891/acta.chem.scand.35a-0727 | url = http://actachemscand.org/pdf/acta_vol_35a_p0727-0728.pdf | doi-access = free }}</ref><ref>{{cite book | author = Vulte, H. T. | title = Laboratory Manual of Inorganic Preparations | publisher = Read Books | year = 2007 | isbn = 978-1-4086-0840-1 }}</ref> |
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the complex {{chem2|Zn(NH3)4Cl2*H2O}} also has been isolated. The latter contains the {{chem2|[Zn(NH3)6](2+)}} ion,<ref name=Wells/> |
the complex {{chem2|Zn(NH3)4Cl2*H2O}} also has been isolated. The latter contains the {{chem2|[Zn(NH3)6](2+)}} ion,.<ref name=Wells/> The species in aqueous solution have been investigated and show that {{chem2|[Zn(NH3)4](2+)}} is the main species present with {{chem2|[Zn(NH3)3Cl]+}} also present at lower {{chem2|NH3}}:Zn ratio.<ref>{{cite journal | last1 = Yamaguchi | first1 = T. | last2 = Ohtaki | first2 = H. | title = X-Ray Diffraction Studies on the Structures of Tetraammine- and Triamminemonochlorozinc(II) Ions in Aqueous Solution | journal = Bulletin of the Chemical Society of Japan | year = 1978 | volume = 51 | issue = 11 | pages = 3227–3231 | doi = 10.1246/bcsj.51.3227 | doi-access = free }}</ref> |
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===Aqueous solutions of zinc chloride=== |
===Aqueous solutions of zinc chloride=== |
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Zinc chloride dissolves readily in water to give {{chem2|ZnCl_{''x''}(H2O)_{4−''x''}|}} species and some free chloride.<ref>{{cite journal | last1 = Irish | first1 = D. E. | last2 = McCarroll | first2 = B. | last3 = Young | first3 = T. F. | title = Raman Study of Zinc Chloride Solutions | journal = The Journal of Chemical Physics | year = 1963 | volume = 39 | issue = 12 | pages = 3436–3444 | doi = 10.1063/1.1734212 | bibcode = 1963JChPh..39.3436I }}</ref><ref>{{cite journal | last1 = Yamaguchi | first1 = T. | last2 = Hayashi | first2 = S. | last3 = Ohtaki | first3 = H. | title = X-Ray Diffraction and Raman Studies of Zinc(II) Chloride Hydrate Melts, ZnCl<sub>2</sub> · ''R'' H<sub>2</sub>O (''R'' = 1.8, 2.5, 3.0, 4.0, and 6.2) | journal = The Journal of Physical Chemistry | year = 1989 | volume = 93 | issue = 6 | pages = 2620–2625 | doi = 10.1021/j100343a074 }}</ref><ref>{{cite journal |author1=Pye, C. C. |author2=Corbeil, C. R. |author3=Rudolph, W. W. | title = An ''ab initio'' Investigation of Zinc Chloro Complexes | journal = Physical Chemistry Chemical Physics | year = 2006 | volume = 8 | issue = 46 | pages = 5428–5436 | doi = 10.1039/b610084h | issn = 1463-9076 | pmid = 17119651 |bibcode=2006PCCP....8.5428P |s2cid=37521287 }}</ref> Aqueous solutions of {{chem2|ZnCl2}} are acidic: a 6 [[concentration#Molarity|M]] aqueous solution has a [[pH]] of 1.<ref name=Holleman/> The acidity of aqueous {{chem2|ZnCl2}} solutions relative to solutions of other Zn<sup>2+</sup> salts (say the sulfate) is due to the formation of the tetrahedral chloro [[aqua complex]]es such as [ZnCl<sub>3</sub>(H<sub>2</sub>O)]<sup>-</sup>.<ref>{{cite book |author = Brown, I. D. | year = 2006 | title = The Chemical Bond in Inorganic Chemistry: The Bond Valence Model | publisher = Oxford University Press | isbn = 978-0-19-929881-5 }}</ref> Most metal dichlorides for octahedral complexes, with stronger O-H bonds. The combination of hydrochloric acid and {{chem2|ZnCl2}} gives a reagent known as "[[Lucas reagent]]". Such reagents were once used a [[qualitative organic analysis|test]] for primary alcohols. Similar reactions are the basis of industrial routes from methanol and ethanol respectively to [[methyl chloride]] and [[ethyl chloride]].<ref>{{cite journal |author1=Kjonaas, R. A. |author2=Riedford, B. A. | title = A Study of the Lucas Test | journal = Journal of Chemical Education | year = 1991 | volume = 68 | issue = 8 | pages = 704 | doi = 10.1021/ed068p704 }}</ref> |
Zinc chloride dissolves readily in water to give {{chem2|ZnCl_{''x''}(H2O)_{4−''x''}|}} species and some free chloride.<ref>{{cite journal | last1 = Irish | first1 = D. E. | last2 = McCarroll | first2 = B. | last3 = Young | first3 = T. F. | title = Raman Study of Zinc Chloride Solutions | journal = The Journal of Chemical Physics | year = 1963 | volume = 39 | issue = 12 | pages = 3436–3444 | doi = 10.1063/1.1734212 | bibcode = 1963JChPh..39.3436I }}</ref><ref>{{cite journal | last1 = Yamaguchi | first1 = T. | last2 = Hayashi | first2 = S. | last3 = Ohtaki | first3 = H. | title = X-Ray Diffraction and Raman Studies of Zinc(II) Chloride Hydrate Melts, ZnCl<sub>2</sub> · ''R'' H<sub>2</sub>O (''R'' = 1.8, 2.5, 3.0, 4.0, and 6.2) | journal = The Journal of Physical Chemistry | year = 1989 | volume = 93 | issue = 6 | pages = 2620–2625 | doi = 10.1021/j100343a074 }}</ref><ref>{{cite journal |author1=Pye, C. C. |author2=Corbeil, C. R. |author3=Rudolph, W. W. | title = An ''ab initio'' Investigation of Zinc Chloro Complexes | journal = Physical Chemistry Chemical Physics | year = 2006 | volume = 8 | issue = 46 | pages = 5428–5436 | doi = 10.1039/b610084h | issn = 1463-9076 | pmid = 17119651 |bibcode=2006PCCP....8.5428P |s2cid=37521287 }}</ref> Aqueous solutions of {{chem2|ZnCl2}} are acidic: a 6 [[concentration#Molarity|M]] aqueous solution has a [[pH]] of 1.<ref name=Holleman/> The acidity of aqueous {{chem2|ZnCl2}} solutions relative to solutions of other Zn<sup>2+</sup> salts (say the sulfate) is due to the formation of the tetrahedral chloro [[aqua complex]]es such as [ZnCl<sub>3</sub>(H<sub>2</sub>O)]<sup>-</sup>.<ref>{{cite book |author = Brown, I. D. | year = 2006 | title = The Chemical Bond in Inorganic Chemistry: The Bond Valence Model | publisher = Oxford University Press | isbn = 978-0-19-929881-5 }}</ref> Most metal dichlorides for octahedral complexes, with stronger O-H bonds. The combination of hydrochloric acid and {{chem2|ZnCl2}} gives a reagent known as "[[Lucas reagent]]". Such reagents were once used a [[qualitative organic analysis|test]] for primary alcohols. Similar reactions are the basis of industrial routes from methanol and ethanol respectively to [[methyl chloride]] and [[ethyl chloride]].<ref>{{cite journal |author1=Kjonaas, R. A. |author2=Riedford, B. A. | title = A Study of the Lucas Test | journal = Journal of Chemical Education | year = 1991 | volume = 68 | issue = 8 | pages = 704 | doi = 10.1021/ed068p704 |bibcode=1991JChEd..68..704K }}</ref> |
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In alkali solution, zinc chloride converts to various zinc hydroxychlorides. These include {{chem2|[Zn(OH)3Cl](2−)}}, {{chem2|[Zn(OH)2Cl2](2−)}}, {{chem2|[Zn(OH)Cl3](2−)}}, and the insoluble {{chem2|Zn5(OH)8Cl2*H2O}}. The latter is the mineral [[simonkolleite]].<ref>{{cite book | author = Zhang, X. G. | year = 1996 | title = Corrosion and Electrochemistry of Zinc | publisher = Springer | isbn = 978-0-306-45334-2 }} {{cite web |url=http://webmineral.com/data/Simonkolleite.shtml#.VEA-9SLF-vM |title= Simonkolleite Mineral Data |author=Staff writer(s)|website= webmineral.com |access-date= October 16, 2014}}</ref> When zinc chloride hydrates are heated, HCl gas evolves and hydroxychlorides result.<ref>{{cite journal |author1=Feigl, F. |author2=Caldas, A. | title = Some Applications of Fusion Reactions with Zinc Chloride in Inorganic Spot Test Analysis | journal = Microchimica Acta | year = 1956 | volume = 44 | issue = 7–8 | pages = 1310–1316 | doi = 10.1007/BF01257465 |s2cid=96823985 }}</ref> |
In alkali solution, zinc chloride converts to various zinc hydroxychlorides. These include {{chem2|[Zn(OH)3Cl](2−)}}, {{chem2|[Zn(OH)2Cl2](2−)}}, {{chem2|[Zn(OH)Cl3](2−)}}, and the insoluble {{chem2|Zn5(OH)8Cl2*H2O}}. The latter is the mineral [[simonkolleite]].<ref>{{cite book | author = Zhang, X. G. | year = 1996 | title = Corrosion and Electrochemistry of Zinc | publisher = Springer | isbn = 978-0-306-45334-2 }} {{cite web |url=http://webmineral.com/data/Simonkolleite.shtml#.VEA-9SLF-vM |title= Simonkolleite Mineral Data |author=Staff writer(s)|website= webmineral.com |access-date= October 16, 2014}}</ref> When zinc chloride hydrates are heated, HCl gas evolves and hydroxychlorides result.<ref>{{cite journal |author1=Feigl, F. |author2=Caldas, A. | title = Some Applications of Fusion Reactions with Zinc Chloride in Inorganic Spot Test Analysis | journal = Microchimica Acta | year = 1956 | volume = 44 | issue = 7–8 | pages = 1310–1316 | doi = 10.1007/BF01257465 |s2cid=96823985 }}</ref> |
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In aqueous solution {{chem2|ZnCl2}}, as well as other halides (bromide, iodide), behave interchangeably for the preparation of other zinc compounds. These salts give |
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⚫ | [[Ninhydrin]] reacts with [[amino acid]]s and [[amine]]s to form a colored compound "Ruhemann's purple" (RP). Spraying with a zinc chloride solution, which is colorless, forms a 1:1 complex RP:{{chem2|ZnCl(H2O)2}}, which is more readily detected as it fluoresces more intensely than RP.<ref>{{cite book | author = Menzel, E. R. | year = 1999 | title = Fingerprint Detection with Lasers | publisher = CRC Press | isbn = 978-0-8247-1974-6 }}</ref> |
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===Redox=== |
===Redox=== |
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Anhydrous zinc chloride melts and even boils without any decomposition up to 900 |
Anhydrous zinc chloride melts and even boils without any decomposition up to 900 °C. These unusual properties invite unusual experiments. One of the very rare examples of zinc compounds that are not Zn<sup>2+</sup>, arise by dissolving zinc metal in molten {{chem2|ZnCl2}} at 500–700 °C. One obtains a yellow diamagnetic solution consisting of the {{chem2|Zn2(2+)}}. The nature of this dimetallic dication has been confirmed by [[Raman spectroscopy]].<ref name="Holleman"/> Although {{chem2|Zn2(2+)}} is unusual, mercury, a heavy congener of zinc, form a wide variety of {{chem2|Hg2(2+)}} salts, see [[mercurous]]. |
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⚫ | In the presence of oxygen, zinc chloride oxidizes to [[zinc oxide]] above 400 °C. Again, this observation indicates the nonoxidation of Zn<sup>2+</sup>.<ref name="decomp">{{cite journal |author1=Frida Jones |author2=Honghi Tran |author3=Daniel Lindberg |author4=Liming Zhao |author5=Mikko Hupa |title=Thermal Stability of Zinc Compounds |journal=Energy & Fuels |date=2013 |volume=27 |issue=10 |pages=5663–5669 |doi=10.1021/ef400505u |language=en}}</ref> |
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===Zinc oxychloride cement=== |
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⚫ | |||
===Zinc hydroxychloride=== |
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⚫ | In the presence of oxygen, zinc chloride oxidizes to [[zinc oxide]] above 400 |
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Concentrated aqueous zinc chloride dissolves [[zinc oxide]] to form zinc hydroxychloride, which is obtained as colorless crystals:<ref>{{cite book|author1=F. Wagenknecht|author2=R. Juza|chapter=Zinc Hydroxychloride|title=Handbook of Preparative Inorganic Chemistry, 2nd Ed. |editor=G. Brauer|publisher=Academic Press|year=1963|place=NY,NY|volume=2pages=1071}}</ref> |
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:{{chem2|ZnCl2 + ZnO + H2O -> 2 ZnCl(OH)}} |
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The same material forms when hydrated zinc chloride is heated.<ref>{{cite book | author = House, J. E. | year = 2008 | title = Inorganic Chemistry | publisher = Academic Press | isbn = 978-0-12-356786-4 }}</ref> |
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: {{chem2|ZnCl2*2H2O → Zn(OH)Cl + HCl + H2O}} |
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===Conversion to other zinc compounds=== |
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Although zinc chlorides with diverse formulas and [[crystal structure|structures]]s exist, these salts often behave very similarly in aqueous solution. For example, solutions prepared from any of the [[water of crystallization|hydrates]] of {{chem2|ZnCl2}}, as well as other halides (bromide, iodide), and the sulfate can often be used interchangeably for the preparation of other zinc compounds. These salts give |
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⚫ | The ability of zinc chloride to dissolve metal oxides (MO)<ref name="HOWI">{{cite book|last=Wiberg|first=Nils|title=Lehrbuch der Anorganischen Chemie |trans-title=Holleman & Wiberg, Textbook of Inorganic chemistry |language=de|publisher=de Gruyter, Berlin|year=2007 |page = 1491|isbn=978-3-11-017770-1}}</ref> is relevant to the utility of {{chem2|ZnCl2}} as a [[flux (metallurgy)|flux]] for [[soldering]]. It dissolves [[Passivation (chemistry)|passivating]] oxides, exposing the clean metal surface.<ref name="HOWI"/><!--Fluxes with {{chem2|ZnCl2}} as an active ingredient are sometimes called "tinner's fluid".{{Citation needed|date=June 2024}} |
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...of the idealized formula {{chem2|MZnOCl2}}.{{Additional citation needed|date=October 2017|reason=Suggested reaction is bit dubious and it is not mentioned in some major chemistry textbooks (or anywhere else that I could find).}}--> |
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==Organic syntheses== |
==Organic syntheses with zinc chloride== |
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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>{{cite journal|title = Hydrocarbons from Methanol|first = Clarence D.|last = Chang|pages = 1–118|doi = 10.1080/01614948308078874|journal = [[Catal. Rev. - Sci. Eng.]]|volume = 25|issue = 1|year = 1983}}</ref> |
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>{{cite journal|title = Hydrocarbons from Methanol|first = Clarence D.|last = Chang|pages = 1–118|doi = 10.1080/01614948308078874|journal = [[Catal. Rev. - Sci. Eng.]]|volume = 25|issue = 1|year = 1983}}</ref> |
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:{{chem2|15 CH3OH → C6(CH3)6 + 3 CH4 + 15 H2O}} |
:{{chem2|15 CH3OH → C6(CH3)6 + 3 CH4 + 15 H2O}} |
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In similar fashion, {{chem2|ZnCl2}} promotes selective [[sodium cyanoborohydride|{{chem2|Na[BH3(CN)]}}]] reduction of tertiary, allylic or benzylic halides to the corresponding hydrocarbons.<ref name="eros" /> |
In similar fashion, {{chem2|ZnCl2}} promotes selective [[sodium cyanoborohydride|{{chem2|Na[BH3(CN)]}}]] reduction of tertiary, allylic or benzylic halides to the corresponding hydrocarbons.<ref name="eros" /> |
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Zinc [[enolate]]s, prepared from alkali metal enolates and {{chem2|ZnCl2}}, 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 {{chem2|ZnCl2}}.<ref>{{cite journal |author1=House, H. O. |author2=Crumrine, D. S. |author3=Teranishi, A. Y. |author4=Olmstead, H. D. | title = Chemistry of Carbanions. XXIII. Use of Metal Complexes to Control the Aldol Condensation | journal = Journal of the American Chemical Society | year = 1973 | volume = 95 | issue = 10 | pages = 3310–3324 | doi = 10.1021/ja00791a039 }}</ref> |
Zinc [[enolate]]s, prepared from alkali metal enolates and {{chem2|ZnCl2}}, 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 {{chem2|ZnCl2}}.<ref>{{cite journal |author1=House, H. O. |author2=Crumrine, D. S. |author3=Teranishi, A. Y. |author4=Olmstead, H. D. | title = Chemistry of Carbanions. XXIII. Use of Metal Complexes to Control the Aldol Condensation | journal = Journal of the American Chemical Society | year = 1973 | volume = 95 | issue = 10 | pages = 3310–3324 | doi = 10.1021/ja00791a039 }}</ref> |
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[[File:ZnCl2 aldol (cropped).gif|center|600px]] |
[[File:ZnCl2 aldol (cropped).gif|center|600px]] |
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==Uses== |
==Uses== |
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===Industrial organic chemistry=== |
===Industrial organic chemistry=== |
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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 in Western countries, |
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>{{cite book |doi=10.1002/14356007.a03_463.pub2 |chapter=Benzaldehyde |title=Ullmann's Encyclopedia of Industrial Chemistry |date=2011 |last1=Brühne |first1=Friedrich |last2=Wright |first2=Elaine |isbn=978-3-527-30385-4 }}</ref> |
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:{{chem2|C6H5CHCl2 + H2O -> C6H5CHO + 2 HCl}} |
:{{chem2|C6H5CHCl2 + H2O -> C6H5CHO + 2 HCl}} |
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Similarly zinc chloride is employed in the main route to [[benzoyl chloride]]. It serves as a catalyst for the production of methylene-bis(dithiocarbamate).<ref name="a">{{Ullmann | author1=Dieter M. M. Rohe | author2=Hans Uwe Wolf | title=Zinc Compounds | year=2007 | pages=1–6 | doi=10.1002/14356007.a28_537|}}</ref> |
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">{{Ullmann | author1=Dieter M. M. Rohe | author2=Hans Uwe Wolf | title=Zinc Compounds | year=2007 | pages=1–6 | doi=10.1002/14356007.a28_537|.pub2}}</ref> |
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===As a metallurgical flux=== |
===As a metallurgical flux=== |
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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. |
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. |
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Zinc chloride forms two salts with ammonium chloride: {{chem2|[NH4]2[ZnCl4]}} and {{chem2|[NH4]3[ZnCl4]Cl}}, 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 {{chem2|H2}} gas and ammonia fumes.<ref>{{cite book | title = ASM handbook | year = 1990 | author = American Society for Metals | publisher = ASM International | isbn = 978-0-87170-021-6 }}</ref> |
Zinc chloride forms two salts with ammonium chloride: {{chem2|[NH4]2[ZnCl4]}} and {{chem2|[NH4]3[ZnCl4]Cl}}, 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 {{chem2|H2}} gas and ammonia fumes.<ref>{{cite book | title = ASM handbook | year = 1990 | author = American Society for Metals | publisher = ASM International | isbn = 978-0-87170-021-6 }}</ref> |
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===In textile and paper processing=== |
===In textile and paper processing=== |
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Relevant to its affinity for these materials, {{chem2|ZnCl2}} is used as a fireproofing agent and in fabric "refresheners" such as Febreze. [[Vulcanized fibre]] is made by soaking paper in concentrated zinc chloride.{{Citation needed|date=June 2024}} |
Relevant to its affinity for these materials, {{chem2|ZnCl2}} is used as a fireproofing agent and in fabric "refresheners" such as Febreze. [[Vulcanized fibre]] is made by soaking paper in concentrated zinc chloride.{{Citation needed|date=June 2024}} |
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===Other uses=== |
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The [[smoke composition|zinc chloride smoke mixture]] ("HC") used in [[smoke grenades]] contains zinc oxide, [[hexachloroethane]] and granular [[aluminium]] powder, which, when ignited, react to form zinc chloride, carbon and [[aluminium oxide]] smoke, an effective [[smoke screen]].<ref>{{cite book | author = Sample, B. E. | year = 1997 | title = Methods for Field Studies of Effects of Military Smokes, Obscurants, and Riot-control Agents on Threatened and Endangered Species | publisher = DIANE Publishing | isbn = 978-1-4289-1233-5 }}</ref> |
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⚫ | [[Ninhydrin]] reacts with [[amino acid]]s and [[amine]]s to form a colored compound "Ruhemann's purple" (RP). Spraying with a zinc chloride solution forms a 1:1 complex RP:{{chem2|ZnCl(H2O)2}}, which is more readily detected as it fluoresces more intensely than RP.<ref>{{cite book | author = Menzel, E. R. | year = 1999 | title = Fingerprint Detection with Lasers | publisher = CRC Press | isbn = 978-0-8247-1974-6 }}</ref> |
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⚫ | Dilute aqueous zinc chloride was used as a disinfectant under the name "Burnett's Disinfecting Fluid".<ref>{{cite book | author = Watts, H. | year = 1869 | title = A Dictionary of Chemistry and the Allied Branches of Other Sciences | publisher = Longmans, Green | url = https://archive.org/details/adictionarychem11wattgoog }}</ref> From 1839 [[Sir William Burnett]] promoted its use as a disinfectant as well as a wood preservative.<ref>{{cite journal |last1=McLean |first1=David |title=Protecting wood and killing germs: 'Burnett's Liquid' and the origins of the preservative and disinfectant industries in early Victorian Britain |journal=Business History |date=April 2010 |volume=52 |issue=2 |pages=285–305|doi=10.1080/00076791003610691 |s2cid=154790730 }}</ref> The Royal Navy conducted trials into its use as a disinfectant in the late 1840s, including during the [[1846–1860 cholera pandemic|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.{{Citation needed|date=June 2024}} |
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==Safety and health== |
==Safety and health== |
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Zinc and chloride are essential for life. Zn<sup>2+</sup> is a component of several [[enzyme]]s, 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>{{cite journal |doi=10.1056/nejmra1810769 |title=Ingestion of Caustic Substances |date=2020 |last1=Hoffman |first1=Robert S. |last2=Burns |first2=Michele M. |last3=Gosselin |first3=Sophie |journal=New England Journal of Medicine |volume=382 |issue=18 |pages=1739–1748 |pmid=32348645 }}</ref> |
Zinc and chloride are essential for life. Zn<sup>2+</sup> is a component of several [[enzyme]]s, 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>{{cite journal |doi=10.1056/nejmra1810769 |title=Ingestion of Caustic Substances |date=2020 |last1=Hoffman |first1=Robert S. |last2=Burns |first2=Michele M. |last3=Gosselin |first3=Sophie |journal=New England Journal of Medicine |volume=382 |issue=18 |pages=1739–1748 |pmid=32348645 }}</ref> Another source of zinc chloride is [[smoke composition|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>{{cite book | author = Sample, B. E. | year = 1997 | title = Methods for Field Studies of Effects of Military Smokes, Obscurants, and Riot-control Agents on Threatened and Endangered Species | publisher = DIANE Publishing | isbn = 978-1-4289-1233-5 }}</ref> Such smoke screens can lead to fatalities.<ref>{{cite book |doi=10.1016/C2011-0-07884-5 |title=Handbook on the Toxicology of Metals |date=2015 |isbn=978-0-444-59453-2|publisher = Academic Press|editor=Gunnar F. Nordberg, Bruce A. Fowler, Monica Nordberg}}</ref> |
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==References== |
==References== |
Latest revision as of 22:50, 20 November 2024
Anhydrous
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Monohydrate
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Names | |
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IUPAC name
Zinc chloride
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Other names
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Identifiers | |
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3D model (JSmol)
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ChEBI | |
ChEMBL | |
ChemSpider | |
DrugBank | |
ECHA InfoCard | 100.028.720 |
EC Number |
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PubChem CID
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RTECS number |
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UNII | |
UN number | 2331 |
CompTox Dashboard (EPA)
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|
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Properties | |
ZnCl2 | |
Molar mass | 136.315 g/mol |
Appearance | White hygroscopic and very deliquescent crystalline solid |
Odor | odorless |
Density | 2.907 g/cm3 |
Melting point | 290 °C (554 °F; 563 K)[1] |
Boiling point | 732 °C (1,350 °F; 1,005 K)[1] |
432.0 g/100 g (25 °C) 615 g/100 g (100 °C) | |
Solubility | soluble in ethanol, glycerol and acetone |
Solubility in ethanol | 430.0 g/100 ml |
−65.0·10−6 cm3/mol | |
Structure | |
Tetrahedral, linear in the gas phase | |
Pharmacology | |
B05XA12 (WHO) | |
Hazards | |
Occupational safety and health (OHS/OSH): | |
Main hazards
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Oral toxicity, irritant[2] |
GHS labelling: | |
Danger | |
H302, H314, H410 | |
P273, P280, P301+P330+P331, P305+P351+P338, P308+P310 | |
NFPA 704 (fire diamond) | |
Lethal dose or concentration (LD, LC): | |
LD50 (median dose)
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|
LC50 (median concentration)
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1260 mg/m3 (rat, 30 min) 1180 mg-min/m3[4] |
NIOSH (US health exposure limits): | |
PEL (Permissible)
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TWA 1 mg/m3 (fume)[3] |
REL (Recommended)
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TWA 1 mg/m3 ST 2 mg/m3 (fume)[3] |
IDLH (Immediate danger)
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50 mg/m3 (fume)[3] |
Safety data sheet (SDS) | External SDS |
Related compounds | |
Other anions
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Other cations
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Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|
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 forms of anhydrous zinc chloride.[5]
All forms of zinc chloride are deliquescent. Zinc chloride finds wide application in textile processing, metallurgical fluxes, and chemical synthesis.[5] In a major monograph, zinc chlorides have been described as "one of the important compounds of zinc."[6]
History
[edit]Zinc chloride has long been known but currently practiced industrial applications all evolved in the latter half of 20th century.[5]
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.[7]
Dilute aqueous zinc chloride was used as a disinfectant under the name "Burnett's Disinfecting Fluid".[8] From 1839 Sir William Burnett promoted its use as a disinfectant as well as a wood preservative.[9] 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.[citation needed]
Structure and properties
[edit]Relative to other metal dihalides, zinc dichloride is unusual in forming several crystalline forms (polymorphs). Four are known: α, β, γ, and δ. Each case features tetrahedral Zn2+ centers.[10]
Form | Crystal system | Pearson symbol | Space group | No. | a (nm) | b (nm) | c (nm) | Z | Density (g/cm3) |
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α | tetragonal | tI12 | I42d | 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.[11][12][13]
The orthorhombic form (δ) rapidly changes to one of the other forms on exposure to the atmosphere. A possible explanation is that the OH− ions originating from the absorbed water facilitate the rearrangement.[10] Rapid cooling of molten ZnCl2 gives a glass.[14]
Molten ZnCl2 has a high viscosity at its melting point and a comparatively low electrical conductivity, which increases markedly with temperature.[15][16] As indicated by a Raman scattering study, the viscosity is explained by the presence of polymers,.[17] Neutron scattering study indicated the presence of tetrahedral ZnCl4 centers, which requires aggregation of ZnCl2 monomers as well.[18]
Hydrates
[edit]Various hydrates of zinc chloride are known: ZnCl2(H2O)n with n = 1, 1.33, 2.5, 3, and 4.5.[19] 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.[20][21]
Preparation and purification
[edit]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[22] The overall method remains useful in industry, but without the solvent:[5]
- Zn + 2 HCl → ZnCl2 + H2
Aqueous solutions may be readily prepared similarly by treating Zn metal, zinc carbonate, zinc oxide, and zinc sulfide with hydrochloric acid:[23]
- ZnS + 2 HCl + 4 H2O → ZnCl2(H2O)4 + H2S
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.[20][24] Lower evaporation temperatures produce higher hydrates.[21]
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.[25] A simple method relies on treating the zinc chloride with thionyl chloride.[26]
Reactions
[edit]Chloride complexes
[edit]A number of salts containing the tetrachlorozincate anion, [ZnCl4]2−, are known.[15] "Caulton's reagent", V2Cl3(thf)6] [Zn2Cl6], which is used in organic chemistry, is an example of a salt containing [Zn2Cl6]2−.[27][28] The compound Cs3ZnCl5 contains tetrahedral [ZnCl4]2− and Cl− anions,[10] so, the compound is not caesium pentachlorozincate, but caesium tetrachlorozincate chloride. No compounds containing the [ZnCl6]4− ion (hexachlorozincate ion) have been characterized.[10] The compound ZnCl2·0.5HCl·H2O crystallizes from a solution of ZnCl2 in hydrochloric acid. It contains a polymeric anion (Zn2Cl−5)n with balancing monohydrated hydronium ions, H5O+2 ions.[10]
Adducts
[edit]The adduct with thf ZnCl2(thf)2 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.[30] A related 1:2 complex is ZnCl2(NH2OH)2 (zinc dichloride di(hydroxylamine)). Known as Crismer's salt, this complexes releases hydroxylamine upon heating.[31] The distinctive ability of aqueous solutions of ZnCl2 to dissolve cellulose is attributed to the formation of zinc-cellulose complexes, illustrating the stability of its adducts.[32] Cellulose also dissolves in molten ZnCl2 hydrate.[33] 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 ZnCl2(NH3)2.[34][35] the complex Zn(NH3)4Cl2·H2O also has been isolated. The latter contains the [Zn(NH3)6]2+ ion,.[10] The species in aqueous solution have been investigated and show that [Zn(NH3)4]2+ is the main species present with [Zn(NH3)3Cl]+ also present at lower NH3:Zn ratio.[36]
Aqueous solutions of zinc chloride
[edit]Zinc chloride dissolves readily in water to give ZnClx(H2O)4−x species and some free chloride.[37][38][39] Aqueous solutions of ZnCl2 are acidic: a 6 M aqueous solution has a pH of 1.[19] The acidity of aqueous ZnCl2 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)]-.[40] Most metal dichlorides for octahedral complexes, with stronger O-H bonds. The combination of hydrochloric acid and ZnCl2 gives a reagent known as "Lucas reagent". Such reagents were once used a test for primary alcohols. Similar reactions are the basis of industrial routes from methanol and ethanol respectively to methyl chloride and ethyl chloride.[41]
In alkali solution, zinc chloride converts to various zinc hydroxychlorides. These include [Zn(OH)3Cl]2−, [Zn(OH)2Cl2]2−, [Zn(OH)Cl3]2−, and the insoluble Zn5(OH)8Cl2·H2O. The latter is the mineral simonkolleite.[42] When zinc chloride hydrates are heated, HCl gas evolves and hydroxychlorides result.[43]
In aqueous solution ZnCl2, 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:[5]
- ZnCl2 + Na2CO3 → ZnCO3 + 2 NaCl
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:ZnCl(H2O)2, which is more readily detected as it fluoresces more intensely than RP.[44]
Redox
[edit]Anhydrous zinc chloride melts and even boils without any decomposition up to 900 °C. These unusual properties invite unusual experiments. One of the very rare examples of zinc compounds that are not Zn2+, arise by dissolving zinc metal in molten ZnCl2 at 500–700 °C. One obtains a yellow diamagnetic solution consisting of the Zn2+2. The nature of this dimetallic dication has been confirmed by Raman spectroscopy.[19] Although Zn2+2 is unusual, mercury, a heavy congener of zinc, form a wide variety of Hg2+2 salts, see mercurous.
In the presence of oxygen, zinc chloride oxidizes to zinc oxide above 400 °C. Again, this observation indicates the nonoxidation of Zn2+.[45]
Zinc hydroxychloride
[edit]Concentrated aqueous zinc chloride dissolves zinc oxide to form zinc hydroxychloride, which is obtained as colorless crystals:[46]
- ZnCl2 + ZnO + H2O → 2 ZnCl(OH)
The same material forms when hydrated zinc chloride is heated.[47]
The ability of zinc chloride to dissolve metal oxides (MO)[48] is relevant to the utility of ZnCl2 as a flux for soldering. It dissolves passivating oxides, exposing the clean metal surface.[48]
Organic syntheses with zinc chloride
[edit]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:[49]
- 15 CH3OH → C6(CH3)6 + 3 CH4 + 15 H2O
This kind of reactivity has been investigated for the valorization of C1 precursors.[50]
Examples of zinc chloride as a Lewis acid include the Fischer indole synthesis:[51]
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.[52] This transformation has in fact been accomplished using even the hydrated ZnCl2 sample shown in the picture above. Many examples describe the use of zinc chloride in Friedel-Crafts acylation reactions.[53][54]
Zinc chloride also activates benzylic and allylic halides towards substitution by weak nucleophiles such as alkenes:[55]
In similar fashion, ZnCl2 promotes selective Na[BH3(CN)] reduction of tertiary, allylic or benzylic halides to the corresponding hydrocarbons.[25]
Zinc enolates, prepared from alkali metal enolates and ZnCl2, 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 ZnCl2.[56]
Organozinc precursor
[edit]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.[57]
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[58] and for the Reformatsky reaction.[59]
Uses
[edit]Industrial organic chemistry
[edit]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:[60]
- C6H5CHCl2 + H2O → C6H5CHO + 2 HCl
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).[5]
As a metallurgical flux
[edit]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: [NH4]2[ZnCl4] and [NH4]3[ZnCl4]Cl, 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.[61]
In textile and paper processing
[edit]Relevant to its affinity for these materials, ZnCl2 is used as a fireproofing agent and in fabric "refresheners" such as Febreze. Vulcanized fibre is made by soaking paper in concentrated zinc chloride.[citation needed]
Safety and health
[edit]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.[5] 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.[62] 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.[63] Such smoke screens can lead to fatalities.[64]
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Further reading
[edit]- N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, 2nd ed., Butterworth-Heinemann, Oxford, UK, 1997.
- Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5.
- 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.