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|Jmol-3D images||Image 1|
|Molar mass||20.01 g mol−1|
|Density||1.15 g/L, gas (25 °C)
0.99 g/mL, liquid (19.5 °C)
-83.6 °C, 190 K, -118 °F
19.5 °C, 293 K, 67 °F
|Solubility in water||miscible|
|Refractive index (nD)||1.00001|
|Dipole moment||1.86 D|
|Std enthalpy of
|−13.66 kJ/g (gas)
−14.99 kJ/g (liquid)
|8.687 J/g K (gas)|
|Other anions||Hydrogen chloride
|Other cations||Sodium fluoride|
|Related compounds||Hydrofluoric acid|
| (what is: /?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
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.
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.
Liquid HF also consists of chains of HF molecules, but the chains are shorter, consisting on average of only five or six molecules. 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.
AcidityThe 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), 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 (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.
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. However, Giguère and Turrell 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−
In concentrated solution, the additional HF causes the ion pair to dissociate with formation of the hydrogen-bonded hydrogen difluoride ion.
- [H3O+•F−] + HF H3O+ + HF2−
- 2 HF H2F+ + F−
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.
Production and usesHydrogen 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.
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.
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.
Acyl chlorides or acid anhydrides react with hydrogen fluoride to give acyl fluorides.
HF is often used in palynology to remove silicate minerals, for extraction of dinoflagellate cysts, acritarchs and chitinozoans.
Health effectshydrofluoric acid, which is highly corrosive and toxic, and requires immediate medical attention upon exposure.
- 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, 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.
- 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 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.
- 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.
- 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.
- F.A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry p. 104
- 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.
- 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
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