Hydrogen fluoride - Misplaced Pages

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)

4 0 1 POI

Flash point

none

Lethal dose or concentration (LD, LC):

LD₅₀ (median dose)

17 ppm (rat, oral)

LC₅₀ (median concentration)

1276 ppm (rat, 1 hr)
1774 ppm (monkey, 1 hr)
4327 ppm (guinea pig, 15 min)

LC_Lo (lowest published)

313 ppm (rabbit, 7 hr)

NIOSH (US health exposure limits):

PEL (Permissible)

TWA 3 ppm

REL (Recommended)

TWA 3 ppm (2.5 mg/m) C 6 ppm (5 mg/m)

IDLH (Immediate danger)

30 ppm

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 , 100 kPa).

Y verify (what is ?) Infobox references

Chemical compound

Hydrogen fluoride (fluorane) is an inorganic compound with chemical formula H F. It is a very poisonous, colorless gas or liquid that dissolves in water to yield hydrofluoric acid. It is the principal industrial source of fluorine, often in the form of hydrofluoric acid, and is an important feedstock in the preparation of many important compounds including pharmaceuticals and polymers such as 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 near room temperature, a much higher temperature than other hydrogen halides.

Hydrogen fluoride is an extremely dangerous gas, forming corrosive and penetrating hydrofluoric acid upon contact with moisture. The gas can also cause blindness by rapid destruction of the corneas.

History

In 1771 Carl Wilhelm Scheele prepared the aqueous solution, hydrofluoric acid in large quantities, although hydrofluoric acid had been known in the 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. Liquid HF also consists of chains of HF molecules, but the chains are shorter, consisting on average of only five or six molecules.

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). 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 HFH₂O with melting point −40 °C (−40 °F), which is 44 °C (79 °F) above the melting point of pure HF.

HF and H₂O similarities
	$graph showing humps of melting temperature, most prominent is at HF 50% mole fraction$
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 hydrofluoric acid. When dilute, hydrofluoric acid behaves like a weak acid, unlike the other hydrohalic acids, due to the formation of hydrogen-bonded ion pairs . However concentrated solutions are strong acids, because bifluoride anions are predominant, instead of ion pairs. In liquid anhydrous HF, self-ionization occurs:

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.

Production

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

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, hydrofluoric acid. Its main uses, on a tonnage basis, are as a precursor to organofluorine compounds and a precursor to cryolite for the electrolysis of aluminium.

Precursor to organofluorine compounds

HF reacts with chlorocarbons to give fluorocarbons. An important application of this reaction is the production of tetrafluoroethylene (TFE), precursor to Teflon. Chloroform is fluorinated by HF to produce chlorodifluoromethane (R-22):

CHCl₃ + 2 HF → CHClF₂ + 2 HCl

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

HF is a reactive solvent in the electrochemical fluorination of organic compounds. In this approach, HF is oxidized in the presence of a hydrocarbon and the fluorine replaces C–H bonds with C–F bonds. Perfluorinated carboxylic acids and sulfonic acids are produced in this way.

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

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.

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.

Catalyst

HF serves as a catalyst in alkylation processes in refineries. It is used in the majority of the installed linear alkyl benzene production facilities in the world. The process involves dehydrogenation of n-paraffins to olefins, and subsequent reaction with benzene using HF as catalyst. For example, in oil refineries "alkylate", a component of high-octane petrol (gasoline), is generated in alkylation units, which combine C₃ and C₄ olefins and iso-butane.

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.

Health effects

Main articles: Hydrofluoric acid and Hydrofluoric acid burn

Hydrogen fluoride is highly corrosive and a powerful contact poison. Exposure requires immediate medical attention. It can cause blindness by rapid destruction of the corneas. Breathing in hydrogen fluoride at high levels or in combination with skin contact can cause death from an irregular heartbeat or from pulmonary edema (fluid buildup in the lungs).

References

^ NIOSH Pocket Guide to Chemical Hazards. "#0334". National Institute for Occupational Safety and Health (NIOSH).
Evans, D. A. "pKa's of Inorganic and Oxo-Acids" (PDF). Retrieved June 19, 2020.
^ "Hydrogen fluoride". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
Johnson, M. W.; Sándor, E.; Arzi, E. (1975). "The Crystal Structure of Deuterium Fluoride". Acta Crystallographica. B31 (8): 1998–2003. doi:10.1107/S0567740875006711.
McLain, Sylvia E.; Benmore, C. J.; Siewenie, J. E.; Urquidi, J.; Turner, J. F. (2004). "On the Structure of Liquid Hydrogen Fluoride". Angewandte Chemie International Edition. 43 (15): 1952–55. doi:10.1002/anie.200353289. PMID 15065271.
Pauling, Linus A. (1960). The Nature of the Chemical Bond and the Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry. Cornell University Press. pp. 454–464. ISBN 978-0-8014-0333-0.
Atkins, Peter; Jones, Loretta (2008). Chemical principles: The quest for insight. W. H. Freeman & Co. pp. 184–185. ISBN 978-1097774678.
Emsley, John (1981). "The hidden strength of hydrogen". New Scientist. 91 (1264): 291–292. Archived from the original on 22 July 2023. Retrieved 25 December 2012.
Greenwood, N. N.; Earnshaw, A. (1998). Chemistry of the Elements (2nd ed.). Oxford: Butterworth Heinemann. pp. 812–816. ISBN 0-7506-3365-4.
C. E. Housecroft and A. G. Sharpe Inorganic Chemistry, p. 221.
F. A. Cotton and G. Wilkinson Advanced Inorganic Chemistry, p. 111.
W. L. Jolly "Modern Inorganic Chemistry" (McGraw-Hill 1984), p. 203. ISBN 0-07-032768-8.
F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry (5th ed.) John Wiley and Sons: New York, 1988. ISBN 0-471-84997-9. p. 109.
^ J. Aigueperse, P. Mollard, D. Devilliers, M. Chemla, R. Faron, R. Romano, J. P. Cuer (2000). "Fluorine Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_307. ISBN 3527306730.{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link)
^ G. Siegemund, W. Schwertfeger, A. Feiring, B. Smart, F. Behr, H. Vogel, B. McKusick (2005). "Fluorine Compounds, Organic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_349. ISBN 978-3527306732.{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link)
M. Jaccaud, R. Faron, D. Devilliers, R. Romano (2005). "Fluorine". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_293. ISBN 978-3527306732.{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link).
Greenwood and Earnshaw, "Chemistry of the Elements", pp. 816–819.
^ Facts About Hydrogen Fluoride (Hydrofluoric Acid)

External links

Fluorides, Hydrogen Fluoride, and Fluorine at ATSDR. Retrieved September 30, 2019
CDC - NIOSH Pocket Guide to Chemical Hazards
Hydrogen Fluoride Fact Sheet at Toxics Use Reduction Institute

v t e Hydrogen compounds
H₃AsO₃ H₃AsO₄ HArF HAt HSO₃F H[BF₄] HBr HBrO HBrO₂ HBrO₃ HBrO₄ HCl HClO HClO₂ HClO₃ HClO₄ HCN HCNO H₂CrO₄/H₂Cr₂O₇ H₂CO₃ H₂CS₃ HF HFO HI HIO HIO₂ HIO₃ HIO₄ HMnO₄ H₂MnO₄ H₂MoO₄ HNC NaHCO₃ HNCO HNO HNO₂ HNO₃ H₂N₂O₂ HNO₅S H₃NSO₃ H₂O H₂O₂ H₂O₃ H₂O₄ H₂O₅ H₃PO₂ H₃PO₃ H₃PO₄ H₄P₂O₇ H₅P₃O₁₀ H₂[PtCl₆] H₂S H₂S₂ H₂Se H₂SeO₃ H₂SeO₄ H₄SiO₄ H₂[SiF₆] HSCN HNCS H₂SO₃ H₂SO₄ H₂SO₅ H₂S₂O₃ H₃O H₂S₂O₆ H₂S₂O₇ H₂S₂O₈ CF₃SO₃H H₂Te H₂TeO₃ H₆TeO₆ H₄TiO₄ H₂Po H[Co(CO)₄]

Hydrogen compounds

Molecules detected in outer space

Molecules

Diatomic	Aluminium monochloride Aluminium monofluoride Aluminium(II) oxide Argonium Carbon cation Carbon monophosphide Carbon monosulfide Carbon monoxide Cyano radical Diatomic carbon Fluoromethylidynium Helium hydride ion Hydrogen chloride Hydrogen fluoride Hydrogen (molecular) Hydroxyl radical Iron(II) oxide Magnesium monohydride Methylidyne radical Nitric oxide Nitrogen (molecular) Imidogen Sulfur mononitride Oxygen (molecular) Phosphorus monoxide Phosphorus mononitride Potassium chloride Silicon carbide Silicon monoxide Silicon monosulfide Sodium chloride Sodium iodide Sulfanyl Sulfur monoxide Titanium(II) oxide
Triatomic	Aluminium(I) hydroxide Aluminium isocyanide Amino radical Carbon dioxide Carbonyl sulfide CCP radical Chloronium Diazenylium Dicarbon monoxide Disilicon carbide Ethynyl radical Formyl radical Hydrogen cyanide (HCN) Hydrogen isocyanide (HNC) Hydrogen sulfide Hydroperoxyl Iron cyanide Isoformyl Magnesium cyanide Magnesium isocyanide Methylene N₂H Nitrous oxide Nitroxyl Ozone Methylidynephosphane Potassium cyanide Trihydrogen cation Sodium cyanide Sodium hydroxide Silicon carbonitride c-Silicon dicarbide SiNC Sulfur dioxide Thioformyl Thioxoethenylidene Titanium dioxide Tricarbon Water
Four atoms	Acetylene Ammonia Isocyanic acid Cyanoethynyl Formaldehyde Fulminic acid HCCN Hydrogen peroxide Hydromagnesium isocyanide Isocyanic acid Isothiocyanic acid Ketenyl Methylene amidogen Methyl cation Methyl radical Propynylidyne Protonated carbon dioxide Protonated hydrogen cyanide Silicon tricarbide Thioformaldehyde Tricarbon monoxide Tricarbon monosulfide Thiocyanic acid
Five atoms	Ammonium ion Butadiynyl Carbodiimide Cyanamide Cyanoacetylene Cyanoformaldehyde Cyanomethyl Cyclopropenylidene Formic acid Isocyanoacetylene Ketene Methane Methoxy radical Methylenimine Propadienylidene Protonated formaldehyde Silane Silicon-carbide cluster
Six atoms	Acetonitrile Cyanobutadiynyl radical E-Cyanomethanimine Cyclopropenone Diacetylene Ethylene Formamide HC₄N Ketenimine Methanethiol Methanol Methyl isocyanide Pentynylidyne Propynal Protonated cyanoacetylene
Seven atoms	Acetaldehyde Acrylonitrile Vinyl cyanide Cyanodiacetylene Ethylene oxide Glycolonitrile Hexatriynyl radical Propyne Methylamine Methyl isocyanate Vinyl alcohol
Eight atoms	Acetic acid Aminoacetonitrile Cyanoallene Ethanimine Glycolaldehyde Hexapentaenylidene Methylcyanoacetylene Methyl formate Acrolein
Nine atoms	Acetamide Cyanohexatriyne Dimethyl ether Ethanol Methyldiacetylene Octatetraynyl radical Propene Ethanethiol Propionitrile N-Methylformamide
Ten atoms or more	Acetone Benzene Benzonitrile Buckminsterfullerene (C₆₀, C₆₀, fullerene, buckyball) C₇₀ fullerene Cyanodecapentayne Ethylene glycol Ethyl formate Methyl acetate Methyl-cyano-diacetylene Methyltriacetylene Propionaldehyde Butyronitrile Pyrimidine Heptatrienyl radical