Ammonium bifluoride

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Ammonium bifluoride
Identifiers
CAS number1341-49-7 YesY
ChemSpider21241205 N
Jmol-3D imagesImage 1
Image 2
Properties
Molecular formulaH5F2N
Molar mass57.04 g mol−1
AppearanceWhite crystals
Density1.50 g cm-3
Melting point

126 °C, 399 K, 259 °F (decomposes)

Boiling point

240 °C, 513 K, 464 °F

Solubility in water63g/100ml 20°C
Solubility in alcoholslightly soluble
Refractive index (nD)1.390
Structure
Crystal structureCubic, related to the CsCl structure
Coordination
geometry
[NH4]+ cation: tetrahedral
[HF2] anion: linear
Hazards
GHS pictogramsGHS-pictogram-acid.svgGHS-pictogram-skull.svg[1]
GHS hazard statementsH301, H314[1]
GHS precautionary statementsP280, P301+310, P305+351+338, P310[1]
NFPA 704
NFPA 704.svg
0
3
0
Related compounds
Other cationspotassium bifluoride
Related compoundsammonium fluoride
 N (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references
 
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Ammonium bifluoride
Identifiers
CAS number1341-49-7 YesY
ChemSpider21241205 N
Jmol-3D imagesImage 1
Image 2
Properties
Molecular formulaH5F2N
Molar mass57.04 g mol−1
AppearanceWhite crystals
Density1.50 g cm-3
Melting point

126 °C, 399 K, 259 °F (decomposes)

Boiling point

240 °C, 513 K, 464 °F

Solubility in water63g/100ml 20°C
Solubility in alcoholslightly soluble
Refractive index (nD)1.390
Structure
Crystal structureCubic, related to the CsCl structure
Coordination
geometry
[NH4]+ cation: tetrahedral
[HF2] anion: linear
Hazards
GHS pictogramsGHS-pictogram-acid.svgGHS-pictogram-skull.svg[1]
GHS hazard statementsH301, H314[1]
GHS precautionary statementsP280, P301+310, P305+351+338, P310[1]
NFPA 704
NFPA 704.svg
0
3
0
Related compounds
Other cationspotassium bifluoride
Related compoundsammonium fluoride
 N (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Ammonium hydrogen fluoride is the inorganic compound with the formula NH4HF2 or NH4F·HF. It is produced from ammonia and hydrogen fluoride. This colourless salt is a glass-etchant and an intermediate in a once-contemplated route to hydrofluoric acid.

Contents

Structure

Ammonium 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.[2] and a bond energy greater than 155 kJ mol−1.[3]

In solid [NH4][HF2], each ammonium cation is surrounded by four fluoride centers in a tetrahedron, with hydrogen - fluorine hydrogen bonds present between the hydrogen atoms of the ammonium ion and the fluorine atoms. Solutions contain tetrahedral [NH4]+ cations and linear [HF2] anions.

Production and applications

Ammonium bifluoride is a component of some etchants. It attacks silica component of glass:

SiO2 + 4 [NH4][HF2] → SiF4 + 4 [NH4]F + 2 H2O

Potassium bifluoride is a related more commonly used etchant.

Ammonium bifluoride has been considered as an intermediate in the production of hydrofluoric acid from hexafluorosilicic acid. Thus, hexafluorosilicic acid is hydrolyzed to give ammonium fluoride, which thermally decomposes to give the bifluoride:

H2SiF6 + 6 NH3 + 2 H2O → SiO2 + 6 NH4F
2 NH4F → NH3 + [NH4]HF2

The resulting ammonium bifluoride is converted to the sodium bifluoride, which thermally decomposes to release HF.[4]

See also

References

  1. ^ a b c Online Sigma Catalogue , accessdate: June 16, 2011.
  2. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth–Heinemann. ISBN 0080379419. 
  3. ^ Emsley, J. (1980) Very strong hydrogen bonds, Chemical Society Reviews, 9, 91-124. doi:10.1039/CS9800900091
  4. ^ Jean Aigueperse, Paul Mollard, Didier Devilliers, Marius Chemla, Robert Faron, Renée Romano, Jean Pierre Cuer (2005), “Fluorine Compounds, Inorganic” in Ullmann’s Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim. doi:10.1002/14356007.a11 307
  1. A. F. Wells (1984) Structural Inorganic Chemistry, 5th ed., Oxford University Press, Oxford, UK.