Hydrogen fluoride

Last updated
Hydrogen fluoride
Hydrogen fluoride.svg
Hydrogen-fluoride-2D-dimensions.svg
Hydrogen-fluoride-3D-vdW.svg
Names
Other names
Fluorane
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.028.759 OOjs UI icon edit-ltr-progressive.svg
KEGG
PubChem CID
RTECS number
  • MW7875000
UNII
UN number 1052
  • InChI=1S/FH/h1H Yes check.svgY
    Key: KRHYYFGTRYWZRS-UHFFFAOYSA-N Yes check.svgY
  • InChI=1/FH/h1H
    Key: KRHYYFGTRYWZRS-UHFFFAOYAC
  • F
Properties
HF
Molar mass 20.006 g·mol−1
Appearancecolourless 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
Std molar
entropy (S⦵298)
8.687 J/g K (gas)
−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:
GHS-pictogram-acid.svg GHS-pictogram-skull.svg GHS-pictogram-exclam.svg
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.svgHealth 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).
Yes check.svgY  verify  (what is  Yes check.svgYX mark.svgN ?)

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 an aqueous solution termed 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, 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.

Contents

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. Hydrogen-fluoride-solid-2D-dimensions.png
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

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
Boiling-points Chalcogen-Halogen.svg HF-H2O Phase-Diagram.svg
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 [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, 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. [14]

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): [14]

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. [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 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. [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

HF burns, not evident until a day after HF burned hands.jpg
HF burns, not evident until a day after

Hydrogen fluoride is highly corrosive and a powerful contact poison. Exposure requires immediate medical attention. [18] 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). [18]

Related Research Articles

<span class="mw-page-title-main">Hydrofluoric acid</span> Solution of hydrogen fluoride in water

Hydrofluoric acid is a solution of hydrogen fluoride (HF) in water. Solutions of HF are colorless, acidic and highly corrosive. It is used to make most fluorine-containing compounds; examples include the commonly used pharmaceutical antidepressant medication fluoxetine (Prozac) and the material PTFE (Teflon). Elemental fluorine is produced from it. It is commonly used to etch glass and silicon wafers.

Calcium fluoride is the inorganic compound of the elements calcium and fluorine with the formula CaF2. It is a white solid that is practically insoluble in water. It occurs as the mineral fluorite (also called fluorspar), which is often deeply coloured owing to impurities.

In chemistry, an interhalogen compound is a molecule which contains two or more different halogen atoms and no atoms of elements from any other group.

<span class="mw-page-title-main">Hydrogen halide</span> Chemical compound consisting of hydrogen bonded to a halogen element

In chemistry, hydrogen halides are diatomic, inorganic compounds that function as Arrhenius acids. The formula is HX where X is one of the halogens: fluorine, chlorine, bromine, iodine, astatine, or tennessine. All known hydrogen halides are gases at Standard Temperature and Pressure.

<span class="mw-page-title-main">Oxygen difluoride</span> Chemical compound

Oxygen difluoride is a chemical compound with the formula OF2. As predicted by VSEPR theory, the molecule adopts a bent molecular geometry. It is a strong oxidizer and has attracted attention in rocketry for this reason. With a boiling point of −144.75 °C, OF2 is the most volatile (isolable) triatomic compound. The compound is one of many known oxygen fluorides.

<span class="mw-page-title-main">Oxygen fluoride</span> Any binary compound of oxygen and fluorine

Oxygen fluorides are compounds of elements oxygen and fluorine with the general formula OnF2, where n = 1 to 6. Many different oxygen fluorides are known:

<span class="mw-page-title-main">Iron(III) fluoride</span> Chemical compound

Iron(III) fluoride, also known as ferric fluoride, are inorganic compounds with the formula FeF3(H2O)x where x = 0 or 3. They are mainly of interest by researchers, unlike the related iron(III) chloride. Anhydrous iron(III) fluoride is white, whereas the hydrated forms are light pink.

Antimony pentafluoride is the inorganic compound with the formula SbF5. This colourless, viscous liquid is a strong Lewis acid and a component of the superacid fluoroantimonic acid, formed upon mixing liquid HF with liquid SbF5 in 1:1 ratio. It is notable for its strong Lewis acidity and the ability to react with almost all known compounds.

<span class="mw-page-title-main">Hexafluorosilicic acid</span> Octahedric silicon compound

Hexafluorosilicic acid is an inorganic compound with the chemical formula H
2SiF
6. Aqueous solutions of hexafluorosilicic acid consist of salts of the cation and hexafluorosilicate anion. These salts and their aqueous solutions are colorless.

<span class="mw-page-title-main">Ammonium bifluoride</span> Chemical compound

Ammonium bifluoride is the inorganic compound with the formula [NH4][HF2] or [NH4]F·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.

<span class="mw-page-title-main">Zirconium tetrafluoride</span> Chemical compound

Zirconium(IV) fluoride describes members of a family inorganic compounds with the formula (ZrF4(H2O)x. All are colorless, diamagnetic solids. Anhydrous Zirconium(IV) fluoride' is a component of ZBLAN fluoride glass.

<span class="mw-page-title-main">Zinc fluoride</span> Chemical compound

Zinc fluoride is an inorganic chemical compound with the chemical formula ZnF2. It is encountered as the anhydrous form and also as the tetrahydrate, ZnF2·4H2O (rhombohedral crystal structure). It has a high melting point and has the rutile structure containing 6 coordinate zinc, which suggests appreciable ionic character in its chemical bonding. Unlike the other zinc halides, ZnCl2, ZnBr2 and ZnI2, it is not very soluble in water.

The bifluoride ion is an inorganic anion with the chemical formula [HF2]−. The anion is colorless. Salts of bifluoride are commonly encountered in the reactions of fluoride salts with hydrofluoric acid. The commercial production of fluorine involves electrolysis of bifluoride salts.

Bromine compounds are compounds containing the element bromine (Br). These compounds usually form the -1, +1, +3 and +5 oxidation states. Bromine is intermediate in reactivity between chlorine and iodine, and is one of the most reactive elements. Bond energies to bromine tend to be lower than those to chlorine but higher than those to iodine, and bromine is a weaker oxidising agent than chlorine but a stronger one than iodine. This can be seen from the standard electrode potentials of the X2/X− couples (F, +2.866 V; Cl, +1.395 V; Br, +1.087 V; I, +0.615 V; At, approximately +0.3 V). Bromination often leads to higher oxidation states than iodination but lower or equal oxidation states to chlorination. Bromine tends to react with compounds including M–M, M–H, or M–C bonds to form M–Br bonds.

<span class="mw-page-title-main">Potassium bifluoride</span> Chemical compound

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

<span class="mw-page-title-main">Potassium heptafluorotantalate</span> Chemical compound

Potassium heptafluorotantalate is an inorganic compound with the formula K2[TaF7]. It is the potassium salt of the heptafluorotantalate anion [TaF7]2−. This white, water-soluble solid is an intermediate in the purification of tantalum from its ores and is the precursor to the metal.

<span class="mw-page-title-main">Thiophosphoryl fluoride</span> Chemical compound

Thiophosphoryl fluoride is an inorganic molecular gas with formula PSF3 containing phosphorus, sulfur and fluorine. It spontaneously ignites in air and burns with a cool flame. The discoverers were able to have flames around their hands without discomfort, and called it "probably one of the coldest flames known". The gas was discovered in 1888.

Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.

<span class="mw-page-title-main">Fluorochemical industry</span> Industry dealing with chemicals from fluorine

The global market for chemicals from fluorine was about US$16 billion per year as of 2006. The industry was predicted to reach 2.6 million metric tons per year by 2015. The largest market is the United States. Western Europe is the second largest. Asia Pacific is the fastest growing region of production. China in particular has experienced significant growth as a fluorochemical market and is becoming a producer of them as well. Fluorite mining was estimated in 2003 to be a $550 million industry, extracting 4.5 million tons per year.

<span class="mw-page-title-main">Sodium bifluoride</span> Chemical compound

Sodium bifluoride is the inorganic compound with the formula Na[HF2]. It is a salt of sodium cation and bifluoride anion. It is a white, water-soluble solid that decomposes upon heating. Sodium bifluoride is non-flammable, hygroscopic, and has a pungent smell. Sodium bifluoride has a number of applications in industry.

References

  1. 1 2 3 4 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. 1 2 "Hydrogen fluoride". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  4. ↑ Johnson, M. W.; Sándor, E.; Arzi, E. (1975). "The Crystal Structure of Deuterium Fluoride". Acta Crystallographica . B31 (8): 1998–2003. doi:10.1107/S0567740875006711.
  5. ↑ 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.
  6. ↑ 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.
  7. ↑ Atkins, Peter; Jones, Loretta (2008). Chemical principles: The quest for insight. W. H. Freeman & Co. pp. 184–185. ISBN   978-1-4292-0965-6.
  8. ↑ Emsley, John (1981). "The hidden strength of hydrogen". New Scientist. 91 (1264): 291–292. Retrieved 25 December 2012.
  9. ↑ Greenwood, N. N.; Earnshaw, A. (1998). Chemistry of the Elements (2nd ed.). Oxford: Butterworth Heinemann. pp. 812–816. 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. ↑ W. L. Jolly "Modern Inorganic Chemistry" (McGraw-Hill 1984), p. 203. ISBN   0-07-032768-8.
  13. ↑ 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.
  14. 1 2 3 4 5 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.{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link)
  15. 1 2 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.{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link)
  16. ↑ 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.{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link).
  17. ↑ Greenwood and Earnshaw, "Chemistry of the Elements", pp. 816–819.
  18. 1 2 Facts About Hydrogen Fluoride (Hydrofluoric Acid)
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