Lithium perchlorate

Lithium perchlorate
IUPAC name
Lithium perchlorate
Other names
Perchloric acid, lithium salt; Lithium Cloricum
7791-03-9 YesY
3D model (Jmol) Interactive image
ChemSpider 133514 YesY
ECHA InfoCard 100.029.307
PubChem 23665649
Molar mass 106.39 g·mol−1
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
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)[1]
Solubility soluble in alcohol, ethyl acetate[1]
Solubility in acetone 137 g/100 g[1]
Solubility in alcohol 1.82 g/g (0 °C, in CH3OH)
1.52 g/g (0 °C, in C2H5OH)
1.05 g/g (25 °C, in C3H7OH)
0.793 g/g (0 °C, in C4H9OH)[1]
105 J/mol·K[1]
125.5 J/mol·K[1]
-380.99 kJ/mol
-254 kJ/mol[1]
Main hazards Oxidizer, irritant
Safety data sheet MSDS
GHS pictograms [2]
GHS signal word Danger
H272, H315, H319, H335[2]
P220, P261, P305+351+338[2]
O Xi
R-phrases R8, R36/37/38
S-phrases S17, S26, S36
NFPA 704
Related compounds
Other anions
Lithium chloride
Lithium hypochlorite
Lithium chlorate
Other cations
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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Infobox references

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.


Inorganic chemistry

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, the latter being over 60% of its mass. It has both the highest oxygen to weight and oxygen to volume ratio of all perchlorates, except beryllium diperchlorate, which is expensive and highly toxic.

Organic chemistry

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.[3]

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.[4]


Lithium perchlorate is also used as an electrolyte in lithium-ion batteries. Lithium perchlorate is chosen over alternative electrolytes 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.[5] 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.[5]


Concentrated solutions of lithium perchlorate (4.5 mol/L) are used as a chaotropic agent to denature proteins.


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.[6]


Perchlorates often give explosive mixtures with organic compounds.[6]


  1. 1 2 3 4 5 6 7
  2. 1 2 3 Sigma-Aldrich Co., Lithium perchlorate. Retrieved on 2014-05-09.
  3. Charette, A. B. "Lithium Perchlorate" in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. doi:10.1002/047084289.
  4. Lithium Perchlorate Product Detail Page
  5. 1 2 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.
  6. 1 2 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
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