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Potassium bifluoride

Potassium bifluoride
Potassium bifluoride.svg
Names
IUPAC name
Potassium bifluoride
Other names
Potassium hydrogen difluoride
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.029.233
RTECS number
  • TS6650000
Properties
HF2K
Molar mass 78.103 g/mol
Appearance colourless solid
Odor slightly acidic
Density 2.37 g/cm3
Melting point 238.7 °C (461.7 °F; 511.8 K)
Boiling point decomposes
24.5 g/100 mL (0 °C)
30.1 g/100mL (10 °C)
39.2 g/100 mL (20 °C)
114.0 g/100 mL (80 °C)
Solubility soluble in ethanol
Structure
monoclinic
Thermochemistry
45.56 J/(mol K) [1]
-417.26 kJ·K−1*mol−1
Hazards
Toxic (T), Corrosive (C)
R-phrases (outdated) R25-34
S-phrases (outdated) S22-26, S37-45
Flash point non flammable
Related compounds
Other anions
Potassium fluoride
Other cations
Sodium bifluoride, ammonium bifluoride
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

Potassium bifluoride is the inorganic compound with the formula KHF2. This colourless salt consists of the potassium cation and the bifluoride (HF2) anion. The salt is used in etchant for glass. Sodium bifluoride is related and is also of commercial use as an etchant as well as in cleaning products.[2]

Nature of the chemical bond in the bifluoride anion

Potassium bifluoride, as its name indicates, contains a bifluoride, or hydrogen(difluoride) anion: HF2. This centrosymmetric triatomic anion features the strongest known hydrogen bond, with a FH length of 114 pm,[3] and a bond energy greater than 155 kJ mol−1.[4]

Synthesis and reactions

The salt was prepared by Edmond Frémy who decomposed it to generate, for the first time, hydrogen fluoride. Potassium bifluoride is prepared by treating potassium carbonate or potassium hydroxide with hydrofluoric acid:

2 HF + KOH → KHF2 + H2O

The electrolysis of KHF2 was used by Henri Moissan to isolate the element fluorine in 1886.

A related material containing two equivalents of HF is also known, KH2F3 (CAS#12178-06-2, m.p. 71.7 C). The industrial production of fluorine entails the electrolysis of molten KH2F3.[2]

See also

References

  1. ^ Westrum, Edgar F., Jr.; Pitzer, Kenneth S. (June 1949). "Thermodynamics of the System KHF2-KF-HF, Including Heat Capacities and Entropies of KHF2, and KF. The Nature of the Hydrogen Bond in KHF2". J. Am. Chem. Soc. 71 (6): 1940–1949. doi:10.1021/ja01174a012.
  2. ^ a b Jean Aigueperse, Paul Mollard, Didier Devilliers, Marius Chemla, Robert Faron, Renée Romano, Jean Pierre Cuer, "Fluorine Compounds, Inorganic" in Ullmann’s Encyclopedia of Industrial Chemistry 2005 Wiley-VCH, Weinheim. doi:10.1002/14356007.a11_307
  3. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  4. ^ Emsley, J. (1980) Very strong hydrogen bonds, Chemical Society Reviews, 9, 91-124.