Potassium bifluoride

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Potassium bifluoride
Potassium bifluoride.svg
IUPAC name
Potassium bifluoride
Other names
Potassium hydrogen difluoride
7789-29-9 N
ChemSpider 35308426 N
Jmol interactive 3D Image
PubChem 11829350
RTECS number TS6650000
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
45.56 J/(mol K) [1]
-417.26 kJ·K−1*mol−1
Toxic (T), Corrosive (C)
R-phrases R25-34
S-phrases 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[edit]

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[edit]

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[edit]


  1. ^ PITZER, KENNETH S.; EDGAR F. WESTRUM, JR (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: 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 0080379419. 
  4. ^ Emsley, J. (1980) Very strong hydrogen bonds, Chemical Society Reviews, 9, 91-124.