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| Names | |||
|---|---|---|---|
| Other names Fluorane | |||
| Identifiers | |||
3D model (JSmol) | |||
| ChEBI | |||
| ChemSpider |
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| ECHA InfoCard | 100.028.759 | ||
| KEGG |
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| RTECS number |
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| UNII | |||
| UN number | 1052 | ||
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| 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) | ||
| 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 | ||
Refractive index (nD) | 1.00001 | ||
| Structure | |||
| Linear | |||
| 1.86D | |||
| 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 | |||
| Occupational safety and health (OHS/OSH): | |||
Main hazards | Highly toxic, corrosive, irritant | ||
| GHS labelling: | |||
    | |||
| 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,P330,P361,P363,P403+P233,P405,P501 | |||
| NFPA 704 (fire diamond) | |||
| 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)[3] | ||
LCLo (lowest published) | 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 | |||
Otheranions | Hydrogen chloride Hydrogen bromide Hydrogen iodide Hydrogen astatide | ||
Othercations | Sodium fluoride Potassium fluoride Rubidium fluoride Caesium fluoride | ||
Related compounds | Water Ammonia | ||
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |||
Hydrogen fluoride (fluorane) is aninorganic compound withchemical formulaHF. It is a very poisonous, colorless gas or liquid that dissolves in water to yieldhydrofluoric acid. It is the principal industrial source offluorine, often in the form of hydrofluoric acid, and is an importantfeedstock in the preparation of many important compounds including pharmaceuticals andpolymers such aspolytetrafluoroethylene (PTFE). HF is also widely used in thepetrochemical industry as a component ofsuperacids. Due to strong and extensivehydrogen bonding, it boils near room temperature, a much higher temperature than otherhydrogen halides.
Hydrogen fluoride is an extremely dangerous gas, formingcorrosive and penetratinghydrofluoric acid upon contact withmoisture. The gas can also causeblindness by rapid destruction of thecorneas.
In 1771Carl Wilhelm Scheele prepared the aqueous solution,hydrofluoric acid, in large quantities, although hydrofluoric acid had been known in theglass industry before then.French chemistEdmond Frémy (1814–1894) is credited with discovering hydrogen fluoride while trying to isolatefluorine.
HF is diatomic in the gas-phase. As a liquid, HF forms relatively stronghydrogen 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]
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 highviscosity in the liquid phase and lower than expected pressure in the gas phase.
HF ismiscible 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 of the hydrogen halides (blue) andhydrogen 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 calledhydrofluoric acid. When dilute, hydrofluoric acid behaves like a weak acid, unlike the other hydrohalic acids, due to the formation of hydrogen-bondedion pairs [H3O+·F−]. However concentrated solutions are strong acids, becausebifluoride anions are predominant, instead of ion pairs. In liquid anhydrous HF,self-ionization occurs:[10][11]
which forms an extremely acidic liquid (H0 = −15.1).
Like water, HF can act as a weak base, reacting withLewis acids to givesuperacids. AHammett acidity function (H0) of −21 is obtained withantimony pentafluoride (SbF5), formingfluoroantimonic acid.[12][13]
Hydrogen fluoride is typically produced by the reaction betweensulfuric acid and pure grades of the mineralfluorite:[14]
About 20% of manufactured HF is a byproduct of fertilizer production, which generateshexafluorosilicic acid. This acid can be degraded to release HF thermally and by hydrolysis:
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 toorganofluorine compounds and a precursor to syntheticcryolite for the electrolysis of aluminium.[14]
HF reacts with chlorocarbons to give fluorocarbons. An important application of this reaction is the production oftetrafluoroethylene (TFE), precursor toTeflon. Chloroform is fluorinated by HF to producechlorodifluoromethane (R-22):[14]
Pyrolysis of chlorodifluoromethane (at 550-750 °C) yields TFE.
HF is a reactive solvent in theelectrochemical fluorination of organic compounds. In this approach, HF is oxidized in the presence of ahydrocarbon and the fluorine replaces C–H bonds withC–F bonds.Perfluorinated carboxylic acids andsulfonic acids are produced in this way.[15]
1,1-Difluoroethane is produced by adding HF toacetylene using mercury as a catalyst.[15]
The intermediate in this process isvinyl fluoride or fluoroethylene, themonomeric precursor topolyvinyl fluoride.
Theelectrowinning ofaluminium 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, includinguranium tetrafluoride.[14]
HF is the precursor to elementalfluorine, F2, byelectrolysis of a solution of HF andpotassium bifluoride. The potassium bifluoride is needed because anhydrous HF does not conduct electricity. Several thousand tons of F2 are produced annually.[16]
HF serves as acatalyst inalkylation processes in refineries. It is used in the majority of the installedlinear alkyl benzene production facilities in the world. The process involves dehydrogenation ofn-paraffins to olefins, and subsequent reaction with benzene using HF as catalyst. For example, inoil refineries "alkylate", a component of high-octane petrol (gasoline), is generated in alkylation units, which combine C3 and C4 olefins andiso-butane.[14]
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]
Hydrogen fluoride is highly corrosive and a powerful contact poison. Exposure requires immediate medical attention.[18] It can cause blindness by rapid destruction of thecorneas. Breathing in hydrogen fluoride at high levels or in combination with skin contact can cause death from anirregular heartbeat or frompulmonary edema (fluid buildup in the lungs).[18] Exposure of the intestinal system to HF solution is known to cause fulminant acute colitis requiring surgical intervention.[19]
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