Hydrogen fluoride
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3D model (JSmol)
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| ChemSpider | |||
| ECHA InfoCard | 100.028.759 | ||
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PubChem CID
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CompTox Dashboard (EPA)
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| Properties | |||
| FH | |||
| Molar mass | 20.006 g·mol−1 | ||
| Appearance | colorless gas | ||
| Density | 1.15 g/L, gas (25 °C) 0.99 g/mL, liquid (19.5 °C) | ||
| Melting point | −83.6 °C (−118.5 °F; 189.6 K) | ||
| Boiling point | 19.5 °C (67.1 °F; 292.6 K) | ||
| miscible | |||
| Acidity (pKa) | 3.2 | ||
Refractive index (nD)
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1.00001 | ||
| Structure | |||
| Linear | |||
| 1.86 D | |||
| Thermochemistry | |||
Std molar
entropy (S⦵298) |
8.687 J/g K (gas) | ||
Std enthalpy of
formation (ΔfH⦵298) |
−13.66 kJ/g (gas) −14.99 kJ/g (liquid) | ||
| Hazards | |||
| NFPA 704 (fire diamond) | |||
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Other anions
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Hydrogen chloride Hydrogen bromide Hydrogen iodide | ||
Other cations
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Sodium fluoride | ||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Hydrogen fluoride is a chemical compound with the formula HF. This colorless gas is the principal industrial source of fluorine, often in the aqueous form as hydrofluoric acid, and thus is the precursor to many important compounds including pharmaceuticals and polymers (e.g. Teflon). HF is widely used in the petrochemical industry and is a component of many superacids. Hydrogen fluoride boils just below room temperature whereas the other hydrogen halides condense at much lower temperatures. Unlike the other hydrogen halides, HF is lighter than air and diffuses relatively quickly through porous substances.
Hydrogen fluoride is a highly dangerous gas, forming corrosive and penetrating hydrofluoric acid upon contact with tissue. The gas can also cause blindness by rapid destruction of the corneas.
Structure
Near or above room temperature, HF is a colorless gas. Below −83.6 °C (−118.5 °F), HF forms orthorhombic crystals, consisting of zig-zag chains of HF molecules. The HF molecules, with a short H–F bond of 95 pm, are linked to neighboring molecules by intermolecular H–F distances of 155 pm.[1]
Liquid HF also consists of chains of HF molecules, but the chains are shorter, consisting on average of only five or six molecules.[2] The higher boiling point of HF relative to analogous species, such as HCl, is attributed to hydrogen bonding between HF molecules, as indicated by the existence of chains even in the liquid state.
Acidity
The acidity of hydrofluoric acid solutions vary with concentration owing to hydrogen-bond interactions of the fluoride ion. Dilute solutions are weakly acidic with an acid ionization constant Ka = 6.6×10−4 (or pKa = 3.18),[3] in contrast to corresponding solutions of the other hydrogen halides which are strong acids. Concentrated solutions of hydrogen fluoride are much more strongly acid than implied by this value, as shown by measurements of the Hammett acidity function H0[4] (or “effective pH”). For 100%, HF has an H0, estimated to be between −10.2 and −11, which is comparable to the value −12 for sulfuric acid.[5][6]
In thermodynamic terms, HF solutions are highly non-ideal, with the activity of HF increasing much more rapidly than its concentration. The weak acidity in dilute solution is sometimes attributed to the high H—F bond strength, which combines with the high dissolution enthalpy of HF to outweigh the more negative enthalpy of hydration of the fluoride ion.[7] However, Giguère and Turrell[8][9] have shown by infrared spectroscopy that the predominant solute species is the hydrogen-bonded ion-pair [H3O+•F−], which suggests that the ionization can be described as a double equilibrium:
- H2O + HF ⇌ [H3O+•F−] ⇌ H3O+ + F−
The first equilibrium to the right and the second to the left, meaning that HF is extensively dissociated, but that the tight ion pairs reduce the thermodynamic activity coefficient of H3O+, so that the solution is effectively less acidic.[10]
In concentrated solution, the additional HF causes the ion pair to dissociate with formation of the hydrogen-bonded hydrogen difluoride ion.[8][10]
- [H3O+•F−] + HF ⇌ H3O+ + HF2−
The increase in free H3O+ due to this reaction accounts for the rapid increase in acidity, while fluoride ions are stabilized (and become less basic) by strong hydrogen bonding to HF to form HF2−. This interaction between the acid and its own conjugate base is an example of homoconjugation. At the limit of 100% liquid HF, there is autoionization
- 2 HF ⇌ H2F+ + F−
that forms an extremely acidic solution (H0 = −11).
The acidity of anhydrous HF can be increased even further by the addition of Lewis acids such as SbF5, which can reduce H0 to −21.[5][6]
Production and uses
Hydrogen fluoride is produced as by the action of sulfuric acid on pure grades of the mineral fluorite and also as a side-product of the extraction of the fertilizer precursor phosphoric acid from various minerals. See also hydrofluoric acid.
The anhydrous compound hydrogen fluoride is more commonly used than its aqueous solution, hydrofluoric acid. HF serves as a catalyst in alkylation processes in oil refineries. A component of high-octane gasoline called "alkylate" is generated in Alkylation units that combine C3 and C4 olefins and isobutane to generate gasoline.[11]
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.[11]
Hydrogen fluoride is an important catalyst used in the majority of the installed linear alkyl benzene production in the world. The process involves dehydrogenation of n-paraffins to olefins, and subsequent reaction with benzene using HF as catalyst.
Elemental fluorine, F2, is prepared by electrolysis of a solution of HF and potassium bifluoride. The potassium bifluoride is needed because anhydrous hydrogen fluoride does not conduct electricity. Several million kilograms of F2 are produced annually.[12]
Acyl chlorides or acid anhydrides react with hydrogen fluoride to give acyl fluorides.[13]
Hydrogen fluoride demand has been declining since early 1990s because of proved chlorinated fluorocarbons harmful effect of stratospheric ozone depletion. The Kyoto Protocol limited fluorocarbons using and therefore stimulated soften in hydrogen fluoride demand. Within 5-year period hydrogen fluoride demand in European Union fell 20%. In 2010 EU-27 hydrogen fluoride consumption amounted 170,455 tons.[14]
Health effects
Upon contact with moisture, including tissue, hydrogen fluoride immediately converts to hydrofluoric acid, which is highly corrosive and toxic, and requires immediate medical attention upon exposure.
References
- ^ Johnson, M. W.; Sándor, E.; Arzi, E. (1975). "The Crystal Structure of Deuterium Fluoride". Acta Crystallographica. B31 (8): 1998–2003. doi:10.1107/S0567740875006711.
{{cite journal}}: CS1 maint: multiple names: authors list (link) - ^ Mclain, Sylvia E.; Benmore, CJ; Siewenie, JE; Urquidi, J; Turner, JF (2004). "On the Structure of Liquid Hydrogen Fluoride". Angewandte Chemie, International Edition. 43 (15): 1952–55. doi:10.1002/anie.200353289. PMID 15065271.
- ^ Ralph H. Petrucci; William S. Harwood; Jeffry D. Madura (2007). General chemistry: principles and modern applications. Pearson/Prentice Hall. p. 691. ISBN 978-0-13-149330-8. Retrieved 22 August 2011.
- ^ H.H. Hyman; et al. (1957). "The Hammett acidity function H0 for HF aqueous solutions". J. Amer. Chem. Soc. 79: 3668. doi:10.1021/ja01571a016.
{{cite journal}}: Explicit use of et al. in:|author=(help) - ^ a b W.L. Jolly “Modern Inorganic Chemistry” (McGraw-Hill 1984), p. 203 ISBN 0070327688
- ^ a b F.A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry John Wiley and Sons: New York, 1988. ISBN 0-471-84997-9 p. 109
- ^ C.E. Housecroft and A.G. Sharpe “Inorganic Chemistry” (Pearson Prentice Hall, 2nd ed. 2005), p. 170.
- ^ a b Giguère, Paul A. and Turrell, Sylvia (1980). "The nature of hydrofluoric acid. A spectroscopic study of the proton-transfer complex H3O+...F−". J. Am. Chem. Soc. 102 (17): 5473. doi:10.1021/ja00537a008.
{{cite journal}}: CS1 maint: multiple names: authors list (link) - ^ Radu Iftimie, Vibin Thomas, Sylvain Plessis, Patrick Marchand, and Patrick Ayotte (2008). "Spectral Signatures and Molecular Origin of Acid Dissociation Intermediates". J. Am. Chem. Soc. 130 (18): 5901. doi:10.1021/ja077846o. PMID 18386892.
{{cite journal}}: CS1 maint: multiple names: authors list (link) - ^ a b F.A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry p. 104
- ^ a b J. Aigueperse, P. Mollard, D. Devilliers, M. Chemla, R. Faron, R. Romano, J. P. Cuer, “Fluorine Compounds, Inorganic” in Ullmann’s Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, 2005 doi:10.1002/14356007.a11_307. Cite error: The named reference "Ullmann" was defined multiple times with different content (see the help page).
- ^ M. Jaccaud, R. Faron, D. Devilliers, R. Romano “Fluorine” in Ullmann’s Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, 2005 doi:10.1002/14356007.a11_293.
- ^ G Olah and S Kuhn "Preparation of Acyl Fluorides with Anhydrous Hydrogen Fluoride. The General Use of the Method of Colson and Fredenhagen" J. Org. Chem., 1961, vol. 26, 237-238.doi:10.1021/jo01060a600
- ^ "EU-27 Hydrogen Fluoride Demand Presents 20% Decrease within Recent 5 Years".
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