Zinc chloride
Zinc chloride hydrate
Names
IUPAC name
Zinc chloride
Other names
  • Butter of zinc
  • Neutral zinc chloride (1:2)
  • Zinc bichloride (archaic)
  • Zinc(II) chloride
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
DrugBank
ECHA InfoCard 100.028.720
EC Number
  • 231-592-0
RTECS number
  • ZH1400000
UNII
UN number 2331
  • InChI=1S/2ClH.Zn/h2*1H;/q;;+2/p-2 checkY
    Key: JIAARYAFYJHUJI-UHFFFAOYSA-L checkY
  • InChI=1/2ClH.Zn/h2*1H;/q;;+2/p-2
    Key: JIAARYAFYJHUJI-NUQVWONBAB
  • Cl[Zn]Cl
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)
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
 Moderately toxic, irritant[2]
GHS labelling:
GHS05: CorrosiveGHS07: Exclamation markGHS09: Environmental hazard
Danger
H302, H314, H410
P273, P280, P301+P330+P331, P305+P351+P338, P308+P310
NFPA 704 (fire diamond)
NFPA 704 four-colored diamond
3
0
0
Lethal dose or concentration (LD, LC):
  • 350 mg/kg (rat, oral)
  • 350 mg/kg (mouse, oral)
  • 200 mg/kg (guinea pig, oral)
  • 1100 mg/kg (rat, oral)
  • 1250 mg/kg (mouse, oral)
[3]
1260 mg/m3 (rat, 30 min)
1180 mg-min/m3[3]
NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 1 mg/m3 (fume)[4]
REL (Recommended)
 TWA 1 mg/m3 ST 2 mg/m3 (fume)[4]
IDLH (Immediate danger)
 50 mg/m3 (fume)[4]
Safety data sheet (SDS) External MSDS
Related compounds
Other anions
Other cations
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)
Infobox references

Zinc chloride is the name of inorganic chemical compounds with the formula ZnCl2. It forms 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] This salt is hygroscopic and even deliquescent. Zinc chloride finds wide application in textile processing, metallurgical fluxes, and chemical synthesis. No mineral with this chemical composition is known aside from the very rare mineral simonkolleite, Zn5(OH)8Cl2·H2O.

Structure and properties

Four crystalline forms (polymorphs) of ZnCl2 are known: α, β, γ, and δ. Each case features tetrahedral Zn2+ centers.[6]

FormCrystal systemPearson symbolSpace groupNo.a (nm)　b (nm)c (nm)ZDensity (g/cm3)
αtetragonaltI12I42d1220.53980.53980.6422343.00
βtetragonaltP6P42/nmc1370.36960.36961.07123.09
γmonoclinicmP36P21/c140.6541.1311.23328122.98
δorthorhombicoP12Pna21330.61250.64430.769342.98

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.[7][8][9]

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.[6] Rapid cooling of molten ZnCl2 gives a glass.[10]

Molten ZnCl2 has a high viscosity at its melting point and a comparatively low electrical conductivity, which increases markedly with temperature.[11][12] As indicated by a Raman scattering study, the viscosity is explained by the presence of polymers,[13]. Neutron scattering study indicated the presence of tetrahedral ZnCl4 centers, which requires aggregation of ZnCl2 monomers as well..[14]

In the gas phase, ZnCl2 molecules are linear with a bond length of 205 pm.

Hydrates

Five hydrates of zinc chloride are known: ZnCl2(H2O)n with n = 1, 1.5, 2.5, 3 and 4.[15] The tetrahydrate ZnCl2(H2O)4 crystallizes from aqueous solutions of zinc chloride.[15]

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, zinc carbonate, zinc oxide, and zinc sulfide with hydrochloric acid:

ZnS + 2 HCl + 4 H2O → ZnCl2(H2O)4 + H2S

Unlike many other elements, zinc essentially exists in only one oxidation state, 2+, which simplifies the 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 sublimate to 400 °C in a stream of dry nitrogen gas.[16] Finally, the simplest method relies on treating the zinc chloride with thionyl chloride.[17]

Reactions

The Zn2+2 ion

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 contains the Zn2+2 ion.[15]

Salts of [ZnCl4]2− and [Zn2Cl6]2− ions

A number of salts containing the tetrachlorozincate anion, [ZnCl4]2−, are known.[11] "Caulton's reagent", V2Cl3(thf)6] [Zn2Cl6], which is used in organic chemistry, is an example of a salt containing [Zn2Cl6]2−.[18][19] The compound Cs3ZnCl5 contains tetrahedral [ZnCl4]2− and Cl− anions,[6] so, the compound is not caesium pentachlorozincate, but caesium tetrachlorozincate chloride. No compounds containing the [ZnCl6]4− ion (hexachlorozincate ion) have been characterized.[6]

Aqueous solutions of zinc chloride

Zinc chloride dissolves readily in water to give ZnClx(H2O)4−x species and some free chloride.[20][21][22] Aqueous solutions of ZnCl2 are acidic: a 6 M aqueous solution has a pH of 1.[15] 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 where the reduction in coordination number from 6 to 4 further reduces the strength of the O–H bonds in the solvated water molecules.[23]

Alkaline solutions of zinc chloride

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.[24] When zinc chloride hydrates are heated, HCl gas evolves and hydroxychlorides result.[25]

Solutions of zinc chloride in ammonia

When solutions of zinc chloride are treated with ammonia, various ammine complexes are produced. These include Zn(NH3)4Cl2·H2O and on concentration ZnCl2(NH3)2.[26] The former contains the [Zn(NH3)6]2+ ion,[6] and the latter is molecular with a distorted tetrahedral geometry.[27] 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.[28]

Zinc oxychloride cement

Aqueous zinc chloride reacts with zinc oxide to form an amorphous cement that was first investigated in 1855 by Stanislas Sorel. Sorel later went on to investigate the related magnesium oxychloride cement, which bears his name.[29]

Zinc hydroxide chloride

When hydrated zinc chloride is heated, one obtains a residue of Zn(OH)Cl e.g.[30]

ZnCl2·2H2O → Zn(OH)Cl + HCl + H2O

Acidified zinc chloride

The compound ZnCl2·0.5HCl·H2O may be prepared by careful precipitation from a solution of ZnCl2 acidified with HCl. It contains a polymeric anion (Zn2Cl−5)n with balancing monohydrated hydronium ions, H5O+2 ions.[6][31]

Cellulose dissolution in aqueous solutions of ZnCl2

Cellulose dissolves in aqueous solutions of ZnCl2, and zinc-cellulose complexes have been detected.[32] Cellulose also dissolves in molten ZnCl2 hydrate and carboxylation and acetylation performed on the cellulose polymer.[33]

Using zinc chloride for preparing other zinc salts

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)

Role in organic chemistry

Zinc chloride is used as a catalyst or reagent in diverse reactions conducted on an industrial scale. The partial hydrolysis of benzal chloride in the presence of zinc chloride is the main route to benzoyl chloride. It serves as a catalyst for the production of methylene-bis(dithiocarbamate).[5]

The combination of hydrochloric acid and ZnCl2, known as the "Lucas reagent", is effective for the preparation of alkyl chlorides from alcohols. Similar reactions are the basis of industrial routes from methanol and ethanol respectively to methyl chloride and ethyl chloride.

Laboratory syntheses

Zinc chloride is a common reagent in the laboratory useful Lewis acid in organic chemistry.[34]

Molten zinc chloride catalyses the conversion of methanol to hexamethylbenzene:[35]

15 CH3OH → C6(CH3)6 + 3 CH4 + 15 H2O

Other examples include catalyzing (A) the Fischer indole synthesis,[36] and also (B) Friedel-Crafts acylation reactions involving activated aromatic rings[37][38]

Related to the latter is the classical preparation of the dye fluorescein from phthalic anhydride and resorcinol, which involves a Friedel-Crafts acylation.[39] This transformation has in fact been accomplished using even the hydrated ZnCl2 sample shown in the picture above.

Zinc chloride also activates benzylic and allylic halides towards substitution by weak nucleophiles such as alkenes:[40]

In similar fashion, ZnCl2 promotes selective Na[BH3(CN)] 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 catalyzed Negishi coupling with aryl halides or vinyl halides.[41] 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.[42] The chelate is more stable when the bulky phenyl group is pseudo-equatorial rather than pseudo-axial, i.e., threo rather than erythro.

Other uses

As a metallurgical flux

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 reacts with metal oxides (MO) to give derivatives of the idealized formula MZnOCl2.[43] This reaction is relevant to the utility of ZnCl2 solution as a flux for soldering — it dissolves passivating oxides, exposing the clean metal surface.[43] Fluxes with ZnCl2 as an active ingredient are sometimes called "tinner's fluid".

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.[44]

In textile and paper processing

Concentrated aqueous solutions of zinc chloride (more than 64% weight/weight zinc chloride in water) are capable of dissolving starch, silk, and cellulose.

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.

Smoke grenades

The 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.[45]

Fingerprint detection

Ninhydrin reacts with amino acids and amines to form a colored compound "Ruhemann's purple" (RP). Spraying with a zinc chloride solution forms a 1:1 complex RP:ZnCl(H2O)2, which is more readily detected as it fluoresces more intensely than RP.[46]

Disinfectant and wood preservative

Dilute aqueous zinc chloride was used as a disinfectant under the name "Burnett's Disinfecting Fluid". [47] From 1839 Sir William Burnett promoted its use as a disinfectant as well as a wood preservative.[48] 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.

Safety

Zinc chloride is a chemical irritant of the eyes, skin, and respiratory system.[5][49]

References

  1. 1 2 O'Neil, M. J.; et al. (2001). The Merck index : an encyclopedia of chemicals, drugs, and biologicals. N. J.: Whitehouse Station. ISBN 978-0911910131.
  2. ↑ Zinc chloride toxicity
  3. 1 2 "Zinc chloride fume". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  4. 1 2 3 NIOSH Pocket Guide to Chemical Hazards. "#0674". National Institute for Occupational Safety and Health (NIOSH).
  5. 1 2 3 Dieter M. M. Rohe; Hans Uwe Wolf (2007). "Zinc Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. pp. 1–6. doi:10.1002/14356007.a28_537. ISBN 978-3527306732.
  6. 1 2 3 4 5 6 Wells, A. F. (1984). Structural Inorganic Chemistry. Oxford: Clarendon Press. ISBN 978-0-19-855370-0.
  7. ↑ Oswald, H. R.; Jaggi, H. (1960). "Zur Struktur der wasserfreien Zinkhalogenide I. Die wasserfreien Zinkchloride". Helvetica Chimica Acta. 43 (1): 72–77. doi:10.1002/hlca.19600430109.
  8. ↑ Brynestad, J.; Yakel, H. L. (1978). "Preparation and Structure of Anhydrous Zinc Chloride". Inorganic Chemistry. 17 (5): 1376–1377. doi:10.1021/ic50183a059.
  9. ↑ Brehler, B. (1961). "Kristallstrukturuntersuchungen an ZnCl2". Zeitschrift für Kristallographie. 115 (5–6): 373–402. Bibcode:1961ZK....115..373B. doi:10.1524/zkri.1961.115.5-6.373.
  10. ↑ Mackenzie, J. D.; Murphy, W. K. (1960). "Structure of Glass-Forming Halides. II. Liquid Zinc Chloride". The Journal of Chemical Physics. 33 (2): 366–369. Bibcode:1960JChPh..33..366M. doi:10.1063/1.1731151.
  11. 1 2 Prince, R. H. (1994). King, R. B. (ed.). Encyclopedia of Inorganic Chemistry. John Wiley & Sons. ISBN 978-0-471-93620-6.
  12. ↑ Ray, H. S. (2006). Introduction to Melts: Molten Salts, Slags and Glasses. Allied Publishers. ISBN 978-81-7764-875-1.
  13. ↑ Danek, V. (2006). Physico-Chemical Analysis of Molten Electrolytes. Elsevier. ISBN 978-0-444-52116-3.
  14. ↑ Price, D. L.; Saboungi, M.-L.; Susman, S.; Volin, K. J.; Wright, A. C. (1991). "Neutron Scattering Function of Vitreous and Molten Zinc Chloride". Journal of Physics: Condensed Matter. 3 (49): 9835–9842. Bibcode:1991JPCM....3.9835P. doi:10.1088/0953-8984/3/49/001. S2CID 250902741.
  15. 1 2 3 4 Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 978-0-12-352651-9.
  16. ↑ Glenn J. McGarvey Jean-François Poisson Sylvain Taillemaud (2016). "Zinc chloride". Encyclopedia of Reagents for Organic Synthesis: 1–20. doi:10.1002/047084289X.rz007.pub3. ISBN 9780470842898.
  17. ↑ Pray, A. P. (1990). Anhydrous Metal Chlorides. Inorganic Syntheses. Vol. 28. pp. 321–322.
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  19. ↑ Bouma, R. J.; Teuben, J. H.; Beukema, W. R.; Bansemer, R. L.; Huffman, J. C.; Caulton, K. G. (1984). "Identification of the Zinc Reduction Product of VCl3 · 3THF as [V2Cl3(THF)6]2[Zn2Cl6]". Inorganic Chemistry. 23 (17): 2715–2718. doi:10.1021/ic00185a033.
  20. ↑ Irish, D. E.; McCarroll, B.; Young, T. F. (1963). "Raman Study of Zinc Chloride Solutions". The Journal of Chemical Physics. 39 (12): 3436–3444. Bibcode:1963JChPh..39.3436I. doi:10.1063/1.1734212.
  21. ↑ Yamaguchi, T.; Hayashi, S.; Ohtaki, H. (1989). "X-Ray Diffraction and Raman Studies of Zinc(II) Chloride Hydrate Melts, ZnCl2 · R H2O (R = 1.8, 2.5, 3.0, 4.0, and 6.2)". The Journal of Physical Chemistry. 93 (6): 2620–2625. doi:10.1021/j100343a074.
  22. ↑ Pye, C. C.; Corbeil, C. R.; Rudolph, W. W. (2006). "An ab initio Investigation of Zinc Chloro Complexes". Physical Chemistry Chemical Physics. 8 (46): 5428–5436. Bibcode:2006PCCP....8.5428P. doi:10.1039/b610084h. ISSN 1463-9076. PMID 17119651. S2CID 37521287.
  23. ↑ Brown, I. D. (2006). The Chemical Bond in Inorganic Chemistry: The Bond Valence Model. Oxford University Press. ISBN 978-0-19-929881-5.
  24. ↑ Zhang, X. G. (1996). Corrosion and Electrochemistry of Zinc. Springer. ISBN 978-0-306-45334-2. Staff writer(s). "Simonkolleite Mineral Data". webmineral.com. Retrieved October 16, 2014.
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  26. ↑ Vulte, H. T. (2007). Laboratory Manual of Inorganic Preparations. Read Books. ISBN 978-1-4086-0840-1.
  27. ↑ Yamaguchi, T.; Lindqvist, O. (1981). "The Crystal Structure of Diamminedichlorozinc(II), ZnCl2(NH3)2. A New Refinement" (PDF). Acta Chemica Scandinavica A. 35 (9): 727–728. doi:10.3891/acta.chem.scand.35a-0727.
  28. ↑ Yamaguchi, T.; Ohtaki, H. (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 (11): 3227–3231. doi:10.1246/bcsj.51.3227.
  29. ↑ Wilson, A. D.; Nicholson, J. W. (1993). Acid-Base Cements: Their Biomedical and Industrial Applications. Cambridge University Press. ISBN 978-0-521-37222-0.
  30. ↑ House, J. E. (2008). Inorganic Chemistry. Academic Press. ISBN 978-0-12-356786-4.
  31. ↑ Mellow, J. W. (1946). A Comprehensive Treatise on Inorganic and Theoretical Chemistry. Longmans, Green.
  32. ↑ Xu, Q.; Chen, L.-F. (1999). "Ultraviolet Spectra and Structure of Zinc-Cellulose Complexes in Zinc Chloride Solution". Journal of Applied Polymer Science. 71 (9): 1441–1446. doi:10.1002/(SICI)1097-4628(19990228)71:9<1441::AID-APP8>3.0.CO;2-G.
  33. ↑ Fischer, S.; Leipner, H.; Thümmler, K.; Brendler, E.; Peters, J. (2003). "Inorganic Molten Salts as Solvents for Cellulose". Cellulose. 10 (3): 227–236. doi:10.1023/A:1025128028462. S2CID 92194004.
  34. ↑ Olah, George A.; Doggweiler, Hans; Felberg, Jeff D.; Frohlich, Stephan; Grdina, Mary Jo; Karpeles, Richard; Keumi, Takashi; Inaba, Shin-ichi; Ip, Wai M.; Lammertsma, Koop; Salem, George; Tabor, Derrick (1984). "Onium Ylide chemistry. 1. Bifunctional acid-base-catalyzed conversion of heterosubstituted methanes into ethylene and derived hydrocarbons. The onium ylide mechanism of the C1 → C2 conversion". J. Am. Chem. Soc. 106 (7): 2143–2149. doi:10.1021/ja00319a039.
  35. ↑ Chang, Clarence D. (1983). "Hydrocarbons from Methanol". Catal. Rev. - Sci. Eng. 25 (1): 1–118. doi:10.1080/01614948308078874.
  36. ↑ Shriner, R. L.; Ashley, W. C.; Welch, E. (1942). "2-Phenylindole". Organic Syntheses. 22: 98. doi:10.15227/orgsyn.022.00981955.; Collective Volume, vol. 3, p. 725
  37. ↑ Cooper, S. R. (1941). "Resacetophenone". Organic Syntheses. 21: 103. doi:10.15227/orgsyn.021.0103.; Collective Volume, vol. 3, p. 761
  38. ↑ Dike, S. Y.; Merchant, J. R.; Sapre, N. Y. (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 (26): 4775–4786. doi:10.1016/S0040-4020(01)86481-4.
  39. ↑ Furnell, B. S. (1989). Vogel's Textbook of Practical Organic Chemistry (5th ed.). New York: Longman/Wiley.
  40. ↑ Bauml, E.; Tschemschlok, K.; Pock, R.; Mayr, H. (1988). "Synthesis of γ-Lactones from Alkenes Employing p-Methoxybenzyl Chloride as +CH2-CO2− Equivalent" (PDF). Tetrahedron Letters. 29 (52): 6925–6926. doi:10.1016/S0040-4039(00)88476-2.
  41. ↑ Kim, S.; Kim, Y. J.; Ahn, K. H. (1983). "Selective Reduction of Tertiary, Allyl, and Benzyl Halides by Zinc-Modified Cyanoborohydride in Diethyl Ether". Tetrahedron Letters. 24 (32): 3369–3372. doi:10.1016/S0040-4039(00)86272-3.
  42. ↑ House, H. O.; Crumrine, D. S.; Teranishi, A. Y.; Olmstead, H. D. (1973). "Chemistry of Carbanions. XXIII. Use of Metal Complexes to Control the Aldol Condensation". Journal of the American Chemical Society. 95 (10): 3310–3324. doi:10.1021/ja00791a039.
  43. 1 2 Wiberg, Nils (2007). Lehrbuch der Anorganischen Chemie [Holleman & Wiberg, Textbook of Inorganic chemistry] (in German). de Gruyter, Berlin. p. 1491. ISBN 978-3-11-017770-1.
  44. ↑ American Society for Metals (1990). ASM handbook. ASM International. ISBN 978-0-87170-021-6.
  45. ↑ Sample, B. E. (1997). Methods for Field Studies of Effects of Military Smokes, Obscurants, and Riot-control Agents on Threatened and Endangered Species. DIANE Publishing. ISBN 978-1-4289-1233-5.
  46. ↑ Menzel, E. R. (1999). Fingerprint Detection with Lasers. CRC Press. ISBN 978-0-8247-1974-6.
  47. ↑ Watts, H. (1869). A Dictionary of Chemistry and the Allied Branches of Other Sciences. Longmans, Green.
  48. ↑ McLean, David (April 2010). "Protecting wood and killing germs: 'Burnett's Liquid' and the origins of the preservative and disinfectant industries in early Victorian Britain". Business History. 52 (2): 285–305. doi:10.1080/00076791003610691. S2CID 154790730.
  49. ↑ "NIOSH Pocket Guide to Chemical Hazards". CDC.gov. Retrieved 30 October 2020.

Further reading

  • 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.
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