Lithium chloride

{{short description|Chemical compound}} {{Chembox | Verifiedfields = changed | Watchedfields = changed | verifiedrevid = 410218986 | ImageFile1 = Lithium chloride.jpg | ImageName1 = Sample of lithium chloride in a watch glass | ImageFile = Lithium-chloride-3D-ionic.png | ImageClass = bg-transparent | ImageSize = 150px | ImageFile2 = File:NaCl polyhedra.svg | ImageClass2 = bg-transparent | ImageCaption2 = __ Li+     __ Cl− | ImageName = Unit cell model of lithium chloride | PIN = Lithium chloride | SystematicName = Lithium(1+) chloride |Section1={{Chembox Identifiers | InChI1 = 1/ClH.Li/h1H;/q;+1/p-1 | InChIKey1 = KWGKDLIKAYFUFQ-REWHXWOFAB | CASNo = 7447-41-8 | CASNo_Ref = {{cascite|correct|CAS}} | ChEMBL_Ref = {{ebicite|changed|EBI}} | ChEMBL = 69710 | PubChem = 433294 | ChemSpiderID = 22449 | ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}} | UNII = G4962QA067 | UNII_Ref = {{fdacite|correct|FDA}} | EINECS = 231-212-3 | UNNumber = 2056 | MeSHName = Lithium+chloride | ChEBI_Ref = {{ebicite|correct|EBI}} | ChEBI = 48607 | RTECS = OJ5950000 | StdInChI_Ref = {{stdinchicite|changed|chemspider}} | StdInChI = 1S/ClH.Li/h1H;/q;+1/p-1 | StdInChIKey_Ref = {{stdinchicite|changed|chemspider}} | StdInChIKey = KWGKDLIKAYFUFQ-UHFFFAOYSA-M | SMILES = [Li+].[Cl-] | InChI = 1S/ClH.Li/h1H;/q;+1/p-1 | InChIKey = KWGKDLIKAYFUFQ-UHFFFAOYSA-M}} |Section2={{Chembox Properties | Formula = LiCl | Li=1|Cl=1 | Appearance = white solid hygroscopic, sharp | Density = 2.068 g/cm3 | Solubility = 68.29 g/100 mL (0 °C) 74.48 g/100 mL (10 °C) 84.25 g/100 mL (25 °C) 88.7 g/100 mL (40 °C) 123.44 g/100 mL (100 °C)<ref name=chemister /> | SolubleOther = soluble in hydrazine, methylformamide, butanol, selenium(IV) oxychloride, 1-propanol<ref name=chemister /> | Solvent1 = methanol | Solubility1 = 45.2 g/100 g (0 °C) 43.8 g/100 g (20 °C) 42.36 g/100 g (25 °C)<ref name=sioc>{{cite book|last1 = Seidell|first1 = Atherton|last2 = Linke|first2 = William F.|year = 1952|title = Solubilities of Inorganic and Organic Compounds|publisher = Van Nostrand|url =https://books.google.com/books?id=k2e5AAAAIAAJ|access-date = 2014-06-02}}</ref> 44.6 g/100 g (60 °C)<ref name=chemister>[http://chemister.ru/Database/properties-en.php?dbid=1&id=614 lithium chloride]</ref> | Solvent2 = ethanol | Solubility2 = 14.42 g/100 g (0 °C) 24.28 g/100 g (20 °C) 25.1 g/100 g (30 °C) 23.46 g/100 g (60 °C)<ref name=sioc /> | Solvent3 = formic acid | Solubility3 = 26.6 g/100 g (18 °C) 27.5 g/100 g (25 °C)<ref name=chemister /> | Solvent4 = acetone | Solubility4 = 1.2 g/100 g (20 °C) 0.83 g/100 g (25 °C) 0.61 g/100 g (50 °C)<ref name=chemister /> | Solvent5 = liquid ammonia | Solubility5 = 0.54 g/100 g (−34 °C)<ref name=chemister /> 3.02 g/100 g (25 °C) | MeltingPtC = 605–614 | BoilingPtC = 1382 | RefractIndex = 1.662 (24 °C) | VaporPressure = 1 torr (785 °C) 10 torr (934 °C) 100 torr (1130 °C)<ref name=chemister /> | Viscosity = 0.87 cP (807 °C)<ref name=chemister /> | MagSus = −24.3·10−6 cm3/mol }} |Section3={{Chembox Structure | MolShape = Linear (gas) | Coordination = Octahedral | Dipole = 7.13 D (gas) }} |Section4={{Chembox Thermochemistry | DeltaHf = −408.27 kJ/mol<ref name=chemister /> | DeltaGf = −384 kJ/mol<ref name=chemister /> | DeltaHc = | Entropy = 59.31 J/mol·K<ref name=chemister /> | HeatCapacity = 48.03 J/mol·K<ref name=chemister /> }} | Section6 = {{Chembox Pharmacology | Pharmacology_ref = | ATCCode_prefix = V04 | ATCCode_suffix = CX11 | ATC_Supplemental = | ATCvet = | Licence_EU = | INN = | INN_EMA = | Licence_US = | Legal_status = | Legal_AU = | Legal_AU_comment = | Legal_CA = | Legal_CA_comment = | Legal_NZ = | Legal_NZ_comment = | Legal_UK = | Legal_UK_comment = | Legal_US = | Legal_US_comment = | Legal_EU = | Legal_EU_comment = | Legal_UN = | Legal_UN_comment = | Pregnancy_category = | Pregnancy_AU = | Pregnancy_AU_comment = | Dependence_liability = | AdminRoutes = | Bioavail = | ProteinBound = | Metabolism = | Metabolites = | OnsetOfAction = | HalfLife = | DurationOfAction = | Excretion = }} |Section7={{Chembox Hazards | ExternalSDS = [http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc07/icsc0711.htm ICSC 0711] | GHSPictograms = {{GHS07}}<ref name="sigma">{{Sigma-Aldrich|id=203637|name=Lithium chloride|accessdate=2014-05-09}}</ref> | GHSSignalWord = Warning | HPhrases = {{H-phrases|302|315|319|335}}<ref name="sigma" /> | PPhrases = {{P-phrases|261|305+351+338}}<ref name="sigma" /> | NFPA-H = 2 | NFPA-R = 0 | NFPA-F = 0 | FlashPt = Non-flammable | LD50 = 526 mg/kg (oral, rat)<ref>[https://pubchem.ncbi.nlm.nih.gov/#tab/sidsrcname=ChemIDplus&query=7447-41-8&input_type=text ChemIDplus - 7447-41-8 - KWGKDLIKAYFUFQ-UHFFFAOYSA-M - Lithium chloride - Similar structures search, synonyms, formulas, resource links, and other chemical information]</ref> }} |Section8={{Chembox Related | OtherAnions = Lithium fluoride Lithium bromide Lithium iodide Lithium astatide | OtherCations = Sodium chloride Potassium chloride Rubidium chloride Caesium chloride Francium chloride }} }}

'''Lithium chloride''' is a chemical compound with the formula Li Cl. The salt is a typical ionic compound (with certain covalent characteristics), although the small size of the Li+ ion gives rise to properties not seen for other alkali metal chlorides, such as extraordinary solubility in polar solvents (83.05 g/100 mL of water at 20 °C) and its hygroscopic properties.<ref name=Ullmann/>

==Chemical properties== thumb|75px|left|Color produced when lithium chloride is heated The salt forms crystalline hydrates, unlike the other alkali metal chlorides.<ref>Holleman, A. F.; Wiberg, E. ''Inorganic Chemistry'' Academic Press: San Diego, 2001. {{ISBN|0-12-352651-5}}.</ref> Mono-, tri-, and pentahydrates are known.<ref>{{cite journal |author1=Hönnerscheid Andreas |author2=Nuss Jürgen |author3=Mühle Claus |author4=Jansen Martin | year = 2003 | title = Die Kristallstrukturen der Monohydrate von Lithiumchlorid und Lithiumbromid | journal = Zeitschrift für anorganische und allgemeine Chemie | volume = 629 | issue = 2| pages = 312–316 | doi = 10.1002/zaac.200390049 |bibcode=2003ZAACh.629..312H }}</ref> The anhydrous salt can be regenerated by heating the hydrates. LiCl also absorbs up to four equivalents of ammonia/mol. As with any other ionic chloride, solutions of lithium chloride can serve as a source of chloride ion, e.g., forming a precipitate upon treatment with silver nitrate: : LiCl + AgNO3 → AgCl + LiNO3

==Preparation== Lithium chloride is produced by treatment of lithium carbonate with hydrochloric acid.<ref name=Ullmann>{{Ullmann |doi=10.1002/14356007.a15_393|title=Lithium and Lithium Compounds|first1=Ulrich |last1=Wietelmann|first2=Richard J.|last2=Bauer|year=2005}}</ref> Anhydrous LiCl is prepared from the hydrate by heating in a stream of hydrogen chloride.

==Uses== ===Commercial applications=== Lithium chloride is mainly used for the production of lithium metal by electrolysis of a LiCl/KCl melt at {{convert|450|C|F}}. LiCl is also used as a brazing flux for aluminium in automobile parts. It is used as a desiccant for drying air streams.<ref name=Ullmann/> In more specialized applications, lithium chloride finds some use in organic synthesis, e.g., as an additive in the Stille reaction. Also, in biochemical applications, it can be used to precipitate RNA from cellular extracts.<ref>{{cite journal | doi= 10.1089/dna.1983.2.329 | author1= Cathala, G. | author2=Savouret, J. | author3=Mendez, B. | author4=West, B. L. | author5=Karin, M. | author6=Martial, J. A. | author7=Baxter, J. D. | title = A Method for Isolation of Intact, Translationally Active Ribonucleic Acid | journal = DNA | year = 1983 | volume = 2 | issue = 4 | pages = 329–335 | pmid = 6198133 }}</ref>

Lithium chloride is also used as a flame colorant to produce dark red flames.

===Niche uses=== Lithium chloride is used as a relative humidity standard in the calibration of hygrometers. At {{convert|25|C|F}} a saturated solution (45.8%) of the salt will yield an equilibrium relative humidity of 11.30%. Additionally, lithium chloride can be used as a hygrometer. This deliquescent salt forms a self-solution when exposed to air. The equilibrium LiCl concentration in the resulting solution is directly related to the relative humidity of the air. The percent relative humidity at {{convert|25|C|F}} can be estimated, with minimal error in the range {{convert|10-30|C}}, from the following first-order equation: RH=107.93-2.11C, where C is solution LiCl concentration, percent by mass.

Molten LiCl is used for the preparation of carbon nanotubes,<ref>{{cite journal | doi=10.1016/j.carbon.2014.05.089 | volume=77 | title=Towards large scale preparation of carbon nanostructures in molten LiCl | journal=Carbon | pages=835–845| year=2014 | last1=Kamali | first1=Ali Reza | last2=Fray | first2=Derek J. | bibcode=2014Carbo..77..835K | doi-access=free }}</ref> graphene<ref>{{cite journal | journal = Nanoscale| year= 2015|volume = 7| issue= 26|pages= 11310–11320 | doi = 10.1039/c5nr01132a| pmid= 26053881| title=Large-scale preparation of graphene by high temperature insertion of hydrogen into graphite|url= https://www.repository.cam.ac.uk/bitstream/1810/248812/1/Kamali%20%26%20Fray%202015%20Nanoscale.pdf| last1= Kamali| first1= Ali Reza| last2= Fray| first2= Derek J.| doi-access= free}}</ref> and lithium niobate.<ref>{{cite journal | doi = 10.1016/j.ceramint.2013.07.085 | volume=40 | title=Preparation of lithium niobate particles via reactive molten salt synthesis method | journal=Ceramics International | pages=1835–1841| year=2014 | last1=Kamali | first1=Ali Reza | last2=Fray | first2=Derek J. }}</ref>

Lithium chloride has been shown to have strong acaricidal properties, being effective against ''Varroa destructor'' in populations of honey bees.<ref>{{Cite journal|last1=Ziegelmann|first1=Bettina|last2=Abele|first2=Elisabeth|date=January 12, 2018|title=Lithium chloride effectively kills the honey bee parasite ''Varroa destructor'' by a systemic mode of action|journal=Scientific Reports|volume=8|issue=1|doi=10.1038/s41598-017-19137-5|pmid=29330449|pmc=5766531|page=683 |bibcode=2018NatSR...8..683Z}}</ref>

Lithium chloride is used as an aversive agent in lab animals to study conditioned place preference and aversion.

==Precautions== Lithium salts affect the central nervous system in a variety of ways. While the citrate, carbonate, and orotate salts are currently used to treat bipolar disorder, other lithium salts including the chloride were used in the past. For a short time in the 1940s lithium chloride was manufactured as a salt substitute for people with hypertension, but this was prohibited after the toxic effects of the compound (tremors, fatigue, nausea) were recognized.<ref name="talbott">{{cite journal | author= Talbott J. H. | title = Use of lithium salts as a substitute for sodium chloride | journal = Arch Intern Med | year = 1950 | volume = 85 | issue = 1 | pages = 1–10 | pmid = 15398859 | doi=10.1001/archinte.1950.00230070023001}}</ref><ref>{{cite journal | author1= L. J. Stone | author2=M. luton | author3=J. Gilroy | title = Lithium Chloride as a Substitute for Sodium Chloride in the Diet | journal = Journal of the American Medical Association | year = 1949 | volume = 139 | issue = 11 | pages = 688–692 | doi =10.1001/jama.1949.02900280004002 | pmid= 18128981 }}</ref><ref>{{cite news | magazine = Time | title = Case of trie Substitute Salt | date = 28 February 1949 | url = http://www.time.com/time/magazine/article/0,9171,799873,00.html| archive-url = https://web.archive.org/web/20070302040542/http://www.time.com/time/magazine/article/0,9171,799873,00.html| url-status = dead| archive-date = March 2, 2007}}</ref> It was, however, noted by J. H. Talbott that many symptoms attributed to lithium chloride toxicity may have also been attributable to sodium chloride deficiency, to the diuretics often administered to patients who were given lithium chloride, or to the patients' underlying conditions.<ref name="talbott" />

==References== {{reflist}} * ''Handbook of Chemistry and Physics'', 71st edition, CRC Press, Ann Arbor, Michigan, 1990. * N. N. Greenwood, A. Earnshaw, ''Chemistry of the Elements'', 2nd ed., Butterworth-Heinemann, Oxford, UK, 1997. * R. Vatassery, titration analysis of LiCl, sat'd in Ethanol by AgNO3 to precipitate AgCl(s). EP of this titration gives %Cl by mass. * H. Nechamkin, ''The Chemistry of the Elements'', McGraw-Hill, New York, 1968.

==External links==

* [https://cdnsciencepub.com/doi/pdf/10.1139/v52-020 Radiochemical measurements of activity coefficients, from Betts & MacKenzie, Can. J. Chem.] {{Commons category|Lithium chloride}}{{Lithium compounds}} {{Chlorides}}

Category:Chlorides Category:Alkali metal chlorides Category:Lithium salts Category:Metal halides Category:Desiccants Category:Rock salt crystal structure