Lithium perchlorate
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| Names | |
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| IUPAC name
Lithium perchlorate
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| Other names
Perchloric acid, lithium salt; Lithium Cloricum
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| Identifiers | |
3D model (JSmol)
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| ChemSpider | |
| ECHA InfoCard | 100.029.307 |
PubChem CID
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| UNII | |
CompTox Dashboard (EPA)
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| Properties | |
| LiClO 4 | |
| Molar mass |
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| Appearance | White crystals |
| Odor | Odorless |
| Density | 2.42 g/cm3 |
| Melting point | 236 °C (457 °F; 509 K) |
| Boiling point | 430 °C (806 °F; 703 K) decomposes from 400 °C |
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| Solubility | Soluble in alcohols, ethyl acetate[1] |
| Solubility in acetone | 137 g/100 g[1] |
| Solubility in alcohols | |
| Solubility in ethyl acetate | 95.2 g/100 g[2] |
| Solubility in ethyl ether | 113.7 g/100 g[2] |
| Structure | |
| Pnma, No. 62 | |
a = 865.7(1) pm, b = 691.29(9) pm, c = 483.23(6) pm[3]
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Formula units (Z)
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4 formula per cell |
| tetrahedral at Cl | |
| Thermochemistry | |
Heat capacity (C)
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105 J/mol·K[1] |
Std molar
entropy (S⦵298) |
125.5 J/mol·K[1] |
Std enthalpy of
formation (ΔfH⦵298) |
−380.99 kJ/mol |
Gibbs free energy (ΔfG⦵)
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−254 kJ/mol[1] |
| Hazards | |
| Occupational safety and health (OHS/OSH): | |
Main hazards
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Oxidizer, irritant |
| GHS labelling: | |
  [4]
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| Danger | |
| H272, H315, H319, H335[4] | |
| P220, P261, P305+P351+P338[4] | |
| NFPA 704 (fire diamond) | |
| Safety data sheet (SDS) | MSDS |
| Related compounds | |
Other anions
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Lithium chloride Lithium hypochlorite Lithium chlorate |
Other cations
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Sodium perchlorate Potassium perchlorate Rubidium perchlorate |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Lithium perchlorate is the inorganic compound with the formula LiClO4. This white or colourless crystalline salt is noteworthy for its high solubility in many solvents. It exists both in anhydrous form and as a trihydrate.
Applications
[edit]Inorganic chemistry
[edit]Lithium perchlorate is used as a source of oxygen in some chemical oxygen generators. It decomposes at about 400 °C, yielding lithium chloride and oxygen:[5]
- LiClO4 → LiCl + 2 O2
Over 60% of the mass of the lithium perchlorate is released as oxygen.[2] 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.[6]
Lithium perchlorate is used as an oxidizer in solid rocket propellants, and to produce red colored flame in pyrotechnic compositions.[2][7]
Organic chemistry
[edit]LiClO4 is highly soluble in organic solvents, even diethyl ether. Such solutions are employed in Diels–Alder reactions, where it is proposed that the Lewis acidic Li+ binds to Lewis basic sites on the dienophile, thereby accelerating the reaction.[8]
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.[9]
Solid lithium perchlorate is found to be a mild and efficient Lewis acid for promoting cyanosilylation of carbonyl compounds under neutral conditions.[10]
Batteries
[edit]Lithium perchlorate is also used as an electrolyte salt in lithium-ion batteries. Lithium perchlorate is chosen over alternative salts such as lithium hexafluorophosphate or lithium tetrafluoroborate when its superior electrical impedance, conductivity, hygroscopicity, and anodic stability properties are of importance to the specific application.[11] However, these beneficial properties are often overshadowed by the electrolyte's strong oxidizing properties, making the electrolyte reactive toward its solvent at high temperatures and/or high current loads. Due to these hazards the battery is often considered unfit for industrial applications.[11]
Biochemistry
[edit]Concentrated solutions of lithium perchlorate (4.5 mol/L) are used as a chaotropic agent to denature proteins.
Production
[edit]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 °C.[12]
Safety
[edit]Perchlorates often give explosive mixtures with organic compounds, finely divided metals, sulfur, and other reducing agents.[12][2]
References
[edit]- ^ a b c d e f g "Lithium perchlorate". chemister.ru.
- ^ a b c d e "Lithium Perchlorate". AMCP 706-187 Military Pyrotechnics - Properties of Materials. US Army Materiel Command. October 1963. pp. 181–182.
- ^ Wickleder, Mathias S. (2003). "Crystal Structure of LiClO4". Zeitschrift für Anorganische und Allgemeine Chemie. 629 (9): 1466–1468. doi:10.1002/zaac.200300114.
- ^ a b c Sigma-Aldrich Co., Lithium perchlorate. Retrieved on 2014-05-09.
- ^ Markowitz, M. M.; Boryta, D. A.; Stewart, Harvey Jr. (1964). "Lithium Perchlorate Oxygen Candle. Pyrochemical Source of Pure Oxygen". Industrial & Engineering Chemistry Product Research and Development. 3 (4): 321–330. doi:10.1021/i360012a016.
- ^ Herbert Ellern (1968). Military and Civilian Pyrotechnics. Chemical Publishing Company. p. 237. ISBN 978-0-8206-0364-3. OL 37082807M.
- ^ Basil T. Fedoroff; Oliver E. Sheffield (January 1975). "Lithium Perchlorate". Encyclopedia of explosives and related items. Vol. 7. Picatinny Arsenal. p. L45. LCCN 61-61759.
- ^ 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.
- ^ [1] Lithium Perchlorate Product Detail Page
- ^ N. Azizi, M.R. Saidi (2003). "An improved synthesis of cyanohydrins in the presence of solid LiClO4 under solvent-free conditions". Journal of Organometallic Chemistry. 688 (1–2): 283–285. doi:10.1016/j.jorganchem.2003.09.014.
- ^ a b Xu, Kang (2004). "Nonaqueous liquid electrolytes for lithium-based rechargeable batteries" (PDF). Chemical Reviews. 104 (10): 4303–4417. doi:10.1021/cr030203g. PMID 15669157. Retrieved 24 February 2014.
- ^ a b 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