Hydrogen fluoride

Source: Wikipedia, the free encyclopedia.
Not to be confused with the element hafnium, symbol Hf.
Hydrogen fluoride
Names
Other names
Fluorane
Identifiers
3D model (
JSmol
)
ChEBI
ChemSpider
ECHA InfoCard
 100.028.759 Edit this at Wikidata
KEGG
RTECS number
  • MW7875000
UNII
UN number 1052
  • InChI=1S/FH/h1H checkY
    Key: KRHYYFGTRYWZRS-UHFFFAOYSA-N checkY
  • InChI=1/FH/h1H
    Key: KRHYYFGTRYWZRS-UHFFFAOYAC
  • F
Properties
HF
Molar mass 20.006 g·mol−1
Appearance colourless gas or colourless liquid (below 19.5 °C)
Odor unpleasant
Density 1.15 g/L, gas (25 °C)
0.99 g/mL, liquid (19.5 °C)
1.663 g/mL, solid (–125 °C)
Melting point −83.6 °C (−118.5 °F; 189.6 K)
Boiling point 19.5 °C (67.1 °F; 292.6 K)
completely miscible (liquid)
Vapor pressure 783 mmHg (20 °C)[1]
Acidity (pKa) 3.17 (in water),

15 (in DMSO) [2]

Conjugate acid
 Fluoronium
Conjugate base
 Fluoride
1.00001
Structure
Linear
1.86 D
Thermochemistry
8.687 J/g K (gas)
Std enthalpy of
formation
(ΔfH⦵298)
 −13.66 kJ/g (gas)
−14.99 kJ/g (liquid)
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Highly toxic, corrosive, irritant
GHS labelling:
GHS05: Corrosive GHS06: ToxicGHS07: Exclamation mark
Danger
H300+H310+H330, H314
P260, P262, P264, P270, P271, P280, P284, P301+P310, P301+P330+P331, P302+P350, P303+P361+P353, P304+P340, P305+P351+P338, P310, P320, P321, P322, P330, P361, P363, P403+P233, P405, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 4: Very short exposure could cause death or major residual injury. E.g. VX gasFlammability 0: Will not burn. E.g. waterInstability 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazard POI: Poisonous
4
0
1
POI
Flash point none
Lethal dose or concentration (LD, LC):
17 ppm (rat, oral)
1276 ppm (rat, 1 hr)
1774 ppm (monkey, 1 hr)
4327 ppm (guinea pig, 15 min)[3]
313 ppm (rabbit, 7 hr)[3]
NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 3 ppm[1]
REL (Recommended)
 TWA 3 ppm (2.5 mg/m3) C 6 ppm (5 mg/m3) [15-minute][1]
IDLH
(Immediate danger)
 30 ppm[1]
Related compounds
Other anions
 Hydrogen chloride
Hydrogen bromide
Hydrogen iodide
Hydrogen astatide
Other cations
 Sodium fluoride
Potassium fluoride
Rubidium fluoride
Caesium fluoride
Related compounds
 Water
Ammonia
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Hydrogen fluoride (fluorane) is an

feedstock in the preparation of many important compounds including pharmaceuticals and polymers, e.g. polytetrafluoroethylene (PTFE). HF is also widely used in the petrochemical industry as a component of superacids. Due to strong and extensive hydrogen bonding, it boils at near room temperature, much higher than other hydrogen halides
.

Hydrogen fluoride is an extremely dangerous gas, forming

blindness by rapid destruction of the corneas
.

History

In 1771

glass industry
before then. French chemist Edmond Frémy (1814–1894) is credited with discovering hydrogen fluoride (HF) while trying to isolate fluorine.

Structure and reactions

The structure of chains of HF in crystalline hydrogen fluoride.

HF is diatomic in the gas-phase. As a liquid, HF forms relatively strong hydrogen bonds, hence its relatively high boiling point. Solid HF consists of zig-zag chains of HF molecules. The HF molecules, with a short covalent H–F bond of 95 pm length, are linked to neighboring molecules by intermolecular H–F distances of 155 pm.[4] Liquid HF also consists of chains of HF molecules, but the chains are shorter, consisting on average of only five or six molecules.[5]

Comparison with other hydrogen halides

Hydrogen fluoride does not boil until 20 °C in contrast to the heavier hydrogen halides, which boil between −85 °C (−120 °F) and −35 °C (−30 °F).[6][7][8] This hydrogen bonding between HF molecules gives rise to high viscosity in the liquid phase and lower than expected pressure in the gas phase.

Aqueous solutions

Main article: hydrofluoric acid

HF is

miscible with water (dissolves in any proportion). In contrast, the other hydrogen halides exhibit limiting solubilities in water. Hydrogen fluoride forms a monohydrate HF.H2O with melting point −40 °C (−40 °F), which is 44 °C (79 °F) above the melting point of pure HF.[9]

HF and H2O similarities
graph showing trend-breaking water and HF boiling points: big jogs up versus a trend that is down with lower molecular weight for the other series members. graph showing humps of melting temperature, most prominent is at HF 50% mole fraction
Boiling points of the hydrogen halides (blue) and hydrogen chalcogenides (red): HF and H2O break trends. Freezing point of HF/ H2O mixtures: arrows indicate compounds in the solid state.

Aqueous solutions of HF are called

ion pairs [H3O+·F−]. However concentrated solutions are strong acids, because bifluoride anions are predominant, instead of ion pairs. In liquid anhydrous HF, self-ionization occurs:[10][11]

3 HF ⇌ H2F+ + HF−2

which forms an extremely acidic liquid (H0 = −15.1).

Reactions with Lewis acids

Like water, HF can act as a weak base, reacting with

Lewis acids to give superacids. A Hammett acidity function (H0) of −21 is obtained with antimony pentafluoride (SbF5), forming fluoroantimonic acid.[12][13]

Production

Hydrogen fluoride is typically produced by the reaction between sulfuric acid and pure grades of the mineral fluorite:[14]

CaF2 + H2SO4 → 2 HF + CaSO4

About 20% of manufactured HF is a byproduct of fertilizer production, which generates hexafluorosilicic acid. This acid can be degraded to release HF thermally and by hydrolysis:

H2SiF6 → 2 HF + SiF4
SiF4 + 2 H2O → 4 HF + SiO2

Use

In general, anhydrous hydrogen fluoride is more common industrially than its aqueous solution,

organofluorine compounds and a precursor to cryolite for the electrolysis of aluminium.[14]

Precursor to organofluorine compounds

HF reacts with chlorocarbons to give fluorocarbons. An important application of this reaction is the production of

Teflon. Chloroform is fluorinated by HF to produce chlorodifluoromethane (R-22):[14]

CHCl3 + 2 HF → CHClF2 + 2 HCl

Pyrolysis of chlorodifluoromethane (at 550-750 °C) yields TFE.

HF is a reactive solvent in the

Perfluorinated carboxylic acids and sulfonic acids are produced in this way.[15]

1,1-Difluoroethane is produced by adding HF to acetylene using mercury as a catalyst.[15]

HC≡CH + 2 HF → CH3CHF2

The intermediate in this process is vinyl fluoride or fluoroethylene, the monomeric precursor to polyvinyl fluoride.

Precursor to metal fluorides and fluorine

The electrowinning of aluminium relies on the electrolysis of aluminium fluoride in molten cryolite. Several kilograms of HF are consumed per ton of Al produced. Other metal fluorides are produced using HF, including uranium tetrafluoride.[14]

HF is the precursor to elemental fluorine, F2, by electrolysis of a solution of HF and potassium bifluoride. The potassium bifluoride is needed because anhydrous HF does not conduct electricity. Several thousand tons of F2 are produced annually.[16]

Catalyst

HF serves as a

iso-butane.[14]

Solvent

Hydrogen fluoride is an excellent solvent. Reflecting the ability of HF to participate in hydrogen bonding, even proteins and carbohydrates dissolve in HF and can be recovered from it. In contrast, most non-fluoride inorganic chemicals react with HF rather than dissolving.[17]

Health effects

Main articles: Hydrofluoric acid and Hydrofluoric acid burn
left and right hands, two views, burned index fingers
HF burns, not evident until a day after

Hydrogen fluoride is highly corrosive and a powerful contact poison. Exposure requires immediate medical attention.

irregular heartbeat or from pulmonary edema (fluid buildup in the lungs).[18]

References

  1. ^ a b c d NIOSH Pocket Guide to Chemical Hazards. "#0334". National Institute for Occupational Safety and Health (NIOSH).
  2. ^ Evans, D. A. "pKa's of Inorganic and Oxo-Acids" (PDF). Retrieved June 19, 2020.
  3. ^ a b "Hydrogen fluoride". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  4. doi:10.1107/S0567740875006711
    .
  5. PMID 15065271
    .
  6. ISBN 978-0-8014-0333-0
    .
  7. ISBN 978-1097774678
    .
  8. ^ Emsley, John (1981). "The hidden strength of hydrogen". New Scientist. 91 (1264): 291–292. Retrieved 25 December 2012.
  9. ISBN 0-7506-3365-4
    .
  10. ^ C. E. Housecroft and A. G. Sharpe Inorganic Chemistry, p. 221.
  11. ^ F. A. Cotton and G. Wilkinson Advanced Inorganic Chemistry, p. 111.
  12. ISBN 0-07-032768-8
    .
  13. ISBN 0-471-84997-9
    . p. 109.
  14. ^
    ISBN 978-3527306732.{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link
    )
  15. ^
    ISBN 978-3527306732.{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link
    )
  16. ISBN 978-3527306732.{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link
    ).
  17. ^ Greenwood and Earnshaw, "Chemistry of the Elements", pp. 816–819.
  18. ^ a b Facts About Hydrogen Fluoride (Hydrofluoric Acid)

External links

Wikimedia Commons has media related to Hydrogen fluoride.
Retrieved from "https://en.wikipedia.org/w/index.php?title=Hydrogen_fluoride&oldid=1212673408"