Lithium tetrafluoroborate
Names
IUPAC name
Lithium tetrafluoroborate
Other names
Borate(1-), tetrafluoro-, lithium
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.034.692
UNII
  • InChI=1S/BF4.Li/c2-1(3,4)5;/q-1;+1 checkY
    Key: UFXJWFBILHTTET-UHFFFAOYSA-N checkY
  • InChI=1/BF4.Li/c2-1(3,4)5;/q-1;+1
    Key: UFXJWFBILHTTET-UHFFFAOYAL
  • [Li+].F[B-](F)(F)F
Properties
LiBF4
Molar mass 93.746 g/mol
Appearance White/grey crystalline solid
Odor odorless
Density 0.852 g/cm3 solid
Melting point 296.5 °C (565.7 °F; 569.6 K)
Boiling point decomposes
Very soluble[1]
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Harmful, causes burns,
hygroscopic.
NFPA 704 (fire diamond)
NFPA 704 four-colored diamond
1
0
1
Safety data sheet (SDS) External MSDS
Related compounds
Other anions
 Tetrafluoroborate,
Related compounds
 Nitrosyl tetrafluoroborate
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 tetrafluoroborate is an inorganic compound with the formula LiBF4. It is a white crystalline powder. It has been extensively tested for use in commercial secondary batteries, an application that exploits its high solubility in nonpolar solvents.[2]

Applications

Although BF4− has high ionic mobility, solutions of its Li+ salt are less conductive than other less associated salts.[2] As an electrolyte in lithium-ion batteries, LiBF4 offers some advantages relative to the more common LiPF6. It exhibits greater thermal stability[3] and moisture tolerance.[4] For example, LiBF4 can tolerate a moisture content up to 620 ppm at room temperature whereas LiPF6 readily hydrolyzes into toxic POF3 and HF gases, often destroying the battery's electrode materials. Disadvantages of the electrolyte include a relatively low conductivity and difficulties forming a stable solid electrolyte interface with graphite electrodes.

Thermal stability

Because LiBF4 and other alkali-metal salts thermally decompose to evolve boron trifluoride, the salt is commonly used as a convenient source of the chemical at the laboratory scale:[5]

LiBF4 → LiF + BF3

Production

LiBF4 is a byproduct in the industrial synthesis of diborane:[5][6]

8 BF3 + 6 LiH → B2H6 + 6 LiBF4

LiBF4 can also be synthesized from LiF and BF3 in an appropriate solvent that is resistant to fluorination by BF3 (e.g. HF, BrF3, or liquified SO2):[5]

LiF + BF3 → LiBF4

References

  1. ↑ GFS-CHEMICALS Archived 2006-03-16 at the Wayback Machine
  2. 1 2 Xu, Kang. "Nonaqueous Liquid Electrolytes for Lithium-Based Rechargeable Batteries."Chemical Reviews 2004, volume 104, pp. 4303-418. doi:10.1021/cr030203g
  3. ↑ S. Zhang; K. Xu; T. Jow (2003). "Low-temperature performance of Li-ion cells with a LiBF4-based electrolyte". Journal of Solid State Electrochemistry. 7 (3): 147–151. doi:10.1007/s10008-002-0300-9. S2CID 96775286. Retrieved 16 February 2014.
  4. ↑ S. S. Zhang; z K. Xu & T. R. Jow (2002). "Study of LiBF4 as an Electrolyte Salt for a Li-Ion Battery". Journal of the Electrochemical Society. 149 (5): A586–A590. Bibcode:2002JElS..149A.586Z. doi:10.1149/1.1466857. Retrieved 16 February 2014.
  5. 1 2 3 Robert Brotherton; Joseph Weber; Clarence Guibert & John Little (2000). "Boron Compounds". Ullmann's Encyclopedia of Industrial Chemistry. p. 10. doi:10.1002/14356007.a04_309. ISBN 3527306730.
  6. ↑ Brauer, Georg (1963). Handbook of Preparative Inorganic Chemistry Vol. 1, 2nd Ed. New York: Academic Press. p. 773. ISBN 978-0121266011.


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