Lithium perchlorate: Difference between revisions

{{short description|Chemical compound}}

| Verifiedfields = changed

| Watchedfields = changed

| verifiedrevid = 451153890

| Name = Lithium perchlorate

| ImageFile = Lithium perchlorate.png

| ImageSize = 150px

| ImageFile1 = Lithiumperchlorat.png

| ImageSize1 = 240px

| ImageCaption1 = <span style="color:#86E074; background-color:#86E074;">__</span> [[lithium|Li]]⁺&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;<span style="color:#31FC02; background-color:#31FC02;">__</span> [[chlorate|Cl]]⁷⁺&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;<span style="color:#FE0300; background-color:#FE0300;">__</span> [[oxygen|O]]²⁻ <br/>Unit cell of lithium perchlorate.

| ImageAlt1 = The orthorhombic unit cell of lithium perchlorate under standard conditions.

| ImageName = Lithium perchlorate

| IUPACName = Lithium perchlorate

| OtherNames = Perchloric acid, lithium salt; Lithium Cloricum

| Section1 = {{Chembox Identifiers

| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}

| CASNo_Ref = {{cascite|correct|CAS}}

| UNII_Ref = {{fdacite|changed|FDA}}

| UNII = Q86SE98C9C

| PubChem = 23665649

| Section2 = {{Chembox Properties

| Formula = {{chem|LiClO|4}}

| MolarMass = {{ubli

| 106.39&nbsp;g/mol (anhydrous)

| 160.44&nbsp;g/mol (trihydrate)

}}

| Appearance = White crystals

| Odor = Odorless

| Density = 2.42&nbsp;g/cm³

| Solubility = {{ubli

| 42.7 g/100 mL (0 °C)

| 49 g/100 mL (10 °C)

| 59.8 g/100 mL (25 °C)

| 71.8 g/100 mL (40 °C)

| 119.5 g/100 mL (80 °C)

| 300 g/100 g (120 °C)<ref name=chemister>{{Cite web | url=http://chemister.ru/Database/properties-en.php?dbid=1&id=612 | title=Lithium perchlorate|website=chemister.ru}}</ref>

}}

| SolubleOther = Soluble in [[alcohols]], [[ethyl acetate]]<ref name=chemister />

| Solubility1 = 137&nbsp;g/100 g<ref name=chemister />

| Solvent1 = acetone

| Solubility2 = {{ubli

| 182&nbsp;g/100 g ([[methanol|CH₃OH]])

| 152&nbsp;g/100 g ([[ethanol|C₂H₅OH]])

| 105&nbsp;g/100 g ([[propanol|C₃H₇OH]])

| 79.3&nbsp;g/100 g ([[1-butanol|n-C₄H₉OH]])

| 58&nbsp;g/100 g ([[isobutanol|i-C₄H₉OH]]<ref name=chemister />

}}

| Solvent2 = alcohols

|Solubility3= 95.2&nbsp;g/100 g<ref name="AMCP">{{cite book|chapter=Lithium Perchlorate|title=AMCP 706-187 Military Pyrotechnics - Properties of Materials|chapter-url=https://archive.org/details/AMCP706187MilitaryPyrotechnicsPropertiesOfMaterials/page/n197|pages=181–182|date=October 1963|publisher=[[US Army Materiel Command]]}}</ref>

|Solvent3=ethyl acetate

|Solubility4=113.7&nbsp;g/100 g<ref name="AMCP"/>

|Solvent4=ethyl ether

| MeltingPtC = 236

| BoilingPtC = 430

| BoilingPt_notes = <br/> decomposes from 400&nbsp;°C

| Section3 = {{Chembox Structure

| UnitCellFormulas = 4 formula per cell

| SpaceGroup = Pnma, No. 62

| LattConst_a = 865.7(1)&nbsp;pm

| LattConst_b = 691.29(9)&nbsp;pm

| LattConst_c = 483.23(6)&nbsp;pm

| LattConst_ref = <ref>{{Cite journal | doi=10.1002/zaac.200300114| title=Crystal Structure of LiClO4| year=2003| last1=Wickleder| first1=Mathias S.| journal=Zeitschrift für Anorganische und Allgemeine Chemie| volume=629| issue=9| pages=1466–1468}}</ref>

| Coordination = [[tetrahedral]] at Cl

| Section4 = {{Chembox Thermochemistry

| DeltaHf = −380.99&nbsp;kJ/mol

| DeltaGf = −254&nbsp;kJ/mol<ref name=chemister />

| Entropy = 125.5&nbsp;J/mol·K<ref name=chemister />

| HeatCapacity = 105&nbsp;J/mol·K<ref name=chemister />

| Section7 = {{Chembox Hazards

| ExternalSDS = [http://www.sciencelab.com/msds.php?msdsId=9924516 MSDS]

| MainHazards = Oxidizer, irritant

| NFPA-H = 2 | NFPA-F = 0 | NFPA-R = 0 | NFPA-S = OX

| GHSPictograms = {{GHS03}}{{GHS07}}<ref name="sigma">{{Sigma-Aldrich|id=431567|name=Lithium perchlorate|accessdate=2014-05-09}}</ref>

| GHSSignalWord = Danger

| HPhrases = {{H-phrases|272|315|319|335}}<ref name="sigma" />

| PPhrases = {{P-phrases|220|261|305+351+338}}<ref name="sigma" />

| Section8 = {{Chembox Related

| OtherAnions = [[Lithium chloride]]<br/>[[Lithium hypochlorite]]<br/>[[Lithium chlorate]]

| OtherCations = [[Sodium perchlorate]]<br/>[[Potassium perchlorate]]<br/>[[Rubidium perchlorate]]

==Applications==

===Inorganic chemistry===

Lithium perchlorate is used as a source of [[oxygen]] in some [[chemical oxygen generator]]s. It decomposes at about 400&nbsp;°C, yielding [[lithium chloride]] and [[oxygen]]:<ref>{{cite journal|title=Lithium Perchlorate Oxygen Candle. Pyrochemical Source of Pure Oxygen|first1=M. M.|last1=Markowitz|first2=D. A.|last2=Boryta|first3=Harvey Jr.|last3=Stewart|journal=Industrial & Engineering Chemistry Product Research and Development|year=1964|volume=3|issue=4|pages=321–330

|doi=10.1021/i360012a016}}</ref>

: LiClO₄ → LiCl + 2 O₂

Over 60% of the mass of the lithium perchlorate is released as oxygen.<ref name="AMCP"/> It has both the highest oxygen to weight and oxygen to volume ratio of all practical [[perchlorate]] salts, and higher oxygen to volume ratio than [[liquid oxygen]].<ref>{{cite book|author=Herbert Ellern|title=Military and Civilian Pyrotechnics|page=237|year=1968|publisher=Chemical Publishing Company|isbn=978-0-8206-0364-3|ol=OL37082807M}}</ref>

Lithium perchlorate is used as an [[oxidizer]] in [[solid rocket propellant]]s, and to produce red [[colored flame]] in pyrotechnic compositions.<ref name="AMCP"/><ref>{{cite book|author1=Basil T. Fedoroff|author2=Oliver E. Sheffield|title=Encyclopedia of explosives and related items|chapter=Lithium Perchlorate|date=January 1975 |chapter-url=https://archive.org/details/DTIC_ADA019502/page/45|volume=7|page=L45|publisher=Picatinny Arsenal|lccn=61-61759}}</ref>

===Organic chemistry===

LiClO₄ is highly soluble in organic solvents, even diethyl ether. Such solutions are employed in [[Diels–Alder reaction|Diels–Alder reactions]], where it is proposed that the [[Lewis acid]]ic Li⁺ binds to Lewis basic sites on the dienophile, thereby accelerating the reaction.<ref>Charette, A. B. "Lithium Perchlorate" in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. {{doi|10.1002/047084289X}}.</ref>

Lithium perchlorate is also used as a [[co-catalyst]] in the coupling of α,β-unsaturated carbonyls with aldehydes, also known as the [[Baylis–Hillman reaction]].<ref>[http://www.sigmaaldrich.com/catalog/search/ProductDetail/sigald/205281] Lithium Perchlorate Product Detail Page</ref>

Solid lithium perchlorate is found to be a mild and efficient Lewis acid for promoting cyanosilylation of carbonyl compounds under neutral conditions.<ref>{{cite journal

| author= N. Azizi, M.R. Saidi

| title = An improved synthesis of cyanohydrins in the presence of solid LiClO4 under solvent-free conditions

| journal = Journal of Organometallic Chemistry

| year = 2003

| volume = 688

| issue = 1–2

| pages = 283–285| doi = 10.1016/j.jorganchem.2003.09.014

}}</ref>

===Batteries===

Lithium perchlorate is also used as an [[electrolyte]] salt in [[lithium-ion battery|lithium-ion batteries]]. Lithium perchlorate is chosen over alternative salts such as [[lithium hexafluorophosphate]] or [[lithium tetrafluoroborate]] when its superior [[electrical impedance]], [[Conductivity (electrolytic)|conductivity]], [[hygroscopicity]], and anodic stability properties are of importance to the specific application.<ref name=Xu>{{cite journal|last=Xu|first=Kang|title=Nonaqueous liquid electrolytes for lithium-based rechargeable batteries|journal=Chemical Reviews|date=2004|volume=104|issue=10|pages=4303–4417|doi=10.1021/cr030203g|pmid=15669157|url=http://is.muni.cz/el/1431/podzim2006/C7780/um/Read/2659326/LiON_ellytes_ChRev04_4303.pdf|access-date=24 February 2014}}</ref> However, these beneficial properties are often overshadowed by the electrolyte's strong [[Oxidizing agent|oxidizing]] properties, making the electrolyte reactive toward its [[solvent]] at high temperatures and/or high [[Ampere|current]] loads. Due to these hazards the battery is often considered unfit for industrial applications.<ref name=Xu />

===Biochemistry===

Concentrated solutions of lithium perchlorate (4.5&nbsp;mol/L) are used as a [[chaotropic agent]] to denature [[protein]]s.

Lithium perchlorate can be manufactured by reaction of [[sodium perchlorate]] with [[lithium chloride]]. It can be also prepared by electrolysis of [[lithium chlorate]] at 200 mA/cm² at temperatures above 20&nbsp;°C.<ref name=Vogt>Helmut Vogt, Jan Balej, John E. Bennett, Peter Wintzer, Saeed Akbar Sheikh, Patrizio Gallone "Chlorine Oxides and Chlorine Oxygen Acids" in Ullmann's Encyclopedia of Industrial Chemistry 2002, Wiley-VCH. {{doi|10.1002/14356007.a06_483}}</ref>

[[Perchlorate]]s often give explosive mixtures with organic compounds, finely divided metals, sulfur, and other reducing agents.<ref name=Vogt/><ref name="AMCP"/>

==External links==

* [http://webbook.nist.gov/cgi/cbook.cgi?ID=C7791039 WebBook page for LiClO4]

{{Perchlorates}}

[[Category:Lithium salts]]

[[Category:Electrolytes]]