electronic drawing hydro fluorife 3d
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| Other names Fluorane | |||
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| ECHA InfoCard | 100.028.759  | ||
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| United nations number | 1052 | ||
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| Properties | |||
| Chemical formula | HF | ||
| Molar mass | 20.006 g·mol−one | ||
| Appearance | colourless gas or colourless liquid (below 19.5 °C) | ||
| Density | 1.15 k/L, gas (25 °C) 0.99 yard/mL, liquid (19.5 °C) one.663 g/mL, solid (–125 °C) | ||
| Melting point | −83.6 °C (−118.five °F; 189.vi K) | ||
| Boiling point | 19.v °C (67.i °F; 292.6 K) | ||
| Solubility in water | completely miscible (liquid) | ||
| Vapor pressure level | 783 mmHg (20 °C)[1] | ||
| Acidity (pK a) | 3.17 (in water), 15 (in DMSO) [ii] | ||
| Cohabit acrid | Fluoronium | ||
| Cohabit base | Fluoride | ||
| Refractive index (n D) | i.00001 | ||
| Structure | |||
| Molecular shape | Linear | ||
| Dipole moment | 1.86 D | ||
| Thermochemistry | |||
| Std molar | 8.687 J/chiliad K (gas) | ||
| Std enthalpy of | −xiii.66 kJ/m (gas) −14.99 kJ/g (liquid) | ||
| Hazards | |||
| GHS labelling: | |||
| Pictograms |   | ||
| Signal word | Danger | ||
| Hazard statements | H300, H310, H314, H330 | ||
| Precautionary statements | 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) | iv 0 i | ||
| Lethal dose or concentration (LD, LC): | |||
| LC50 (median concentration) | 1276 ppm (rat, 1 hr) 1774 ppm (monkey, 1 hr) 4327 ppm (guinea sus scrofa, fifteen min)[3] | ||
| LCLo (everyman published) | 313 ppm (rabbit, 7 hr)[3] | ||
| NIOSH (U.s.a. health exposure limits): | |||
| PEL (Permissible) | TWA 3 ppm[ane] | ||
| REL (Recommended) | TWA three ppm (ii.5 mg/mthree) C six ppm (5 mg/mthree) [xv-minute][ane] | ||
| IDLH (Immediate danger) | 30 ppm[i] | ||
| 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). | |||
 Yverify (what is  Y  N ?) | |||
| Infobox references | |||
Hydrogen fluoride is a chemical compound with the chemic formula HF. This colorless gas or liquid is the principal industrial source of fluorine, oftentimes as an aqueous solution called hydrofluoric acrid. It is an important feedstock in the preparation of many important compounds including pharmaceuticals and polymers, e.k. polytetrafluoroethylene (PTFE). HF is widely used in the petrochemical industry as a component of superacids. Hydrogen fluoride boils at well-nigh room temperature, much higher than other hydrogen halides.
Hydrogen fluoride is a highly unsafe gas, forming corrosive and penetrating hydrofluoric acrid upon contact with moisture. The gas tin too cause incomprehension by rapid destruction of the corneas.
History [edit]
In 1771 Carl Wilhelm Scheele prepared the aqueous solution, hydrofluoric acrid in large quantities, although hydrofluoric acrid had been known in the drinking glass industry before then. French chemist Edmond Frémy (1814–1894) is credited with discovering anhydrous hydrogen fluoride (HF) while trying to isolate fluorine.
Structure and reactions [edit]
Although a diatomic molecule, HF forms relatively strong intermolecular hydrogen bonds. Solid HF consists of zig-zag chains of HF molecules. The HF molecules, with a short H–F bail of 95 pm, are linked to neighboring molecules by intermolecular H–F distances of 155 pm.[4] Liquid HF besides consists of chains of HF molecules, merely the chains are shorter, consisting on boilerplate of just 5 or six molecules.[5]
Comparison with other hydrogen halides [edit]
Hydrogen fluoride does not boil until 20 °C in dissimilarity to the heavier hydrogen halides, which boil between −85 °C (−120 °F) and −35 °C (−thirty °F).[6] [7] [8] This hydrogen bonding between HF molecules gives rising to high viscosity in the liquid phase and lower than expected pressure in the gas phase.
Aqueous solutions [edit]
HF is miscible with water (dissolves in any proportion). In contrast, the other hydrogen halides exhibit limiting solubilities in h2o. Hydrogen fluoride forms a monohydrate HF.H2O with chiliad.p.−40 °C (−40 °F), which is 44 °C (79 °F) in a higher place the melting indicate of pure HF.[9]
Aqueous solutions of HF are called hydrofluoric acid. When dilute, hydrofluoric acid behaves like a weak acrid, unlike the other hydrohalic acids, due to the germination 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] [eleven]
- 3 HF ⇌ HiiF+ + HF −
2
which forms an extremely acidic liquid ( H 0 = −15.one).
Reactions with Lewis acids [edit]
Like water, HF tin can human activity as a weak base of operations, reacting with Lewis acids to requite superacids. A Hammett acerbity part (H 0) of −21 is obtained with antimony pentafluoride (SbF5), forming fluoroantimonic acid.[12] [13]
Product [edit]
Hydrogen fluoride is produced past the action of sulfuric acrid on pure grades of the mineral fluorite:[14]
- CaFtwo + H2SOfour → 2 HF + CaSO4 The reaction is endothermic.
Virtually 20% of manufactured HF is a byproduct of fertilizer product, which generates hexafluorosilicic acrid. This acid can exist degraded to release HF thermally and by hydrolysis:
- H2SiFvi → 2 HF + SiF4
- SiFiv + two HiiO → four HF + SiO2
Use [edit]
In general, the anhydrous compound hydrogen fluoride is more common industrially than its aqueous solution, hydrofluoric acid. Its main uses, on a tonnage basis, are as a forerunner to organofluorine compounds and a forerunner to cryolite for the electrolysis of aluminium.[fourteen]
Precursor to organofluorine compounds [edit]
HF reacts with chlorocarbons to give fluorocarbons. An important application of this reaction is the product 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 manner.[15]
1,one-Difluoroethane is produced by adding HF to acetylene using mercury every bit a catalyst.[fifteen]
- HC≡CH + ii HF → CH3CHF2
The intermediate in this process is vinyl fluoride or fluoroethylene, the monomeric precursor to polyvinyl fluoride.
Forerunner to metal fluorides and fluorine [edit]
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 hexafluoride.[fourteen]
HF is the forerunner 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]
Goad [edit]
HF serves as a goad in alkylation processes in refineries. It is used in the majority of the installed linear alkyl benzene production facilities in the globe. The process involves dehydrogenation of n-paraffins to olefins, and subsequent reaction with benzene using HF every bit 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.[fourteen]
Solvent [edit]
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 information technology. In dissimilarity, most non-fluoride inorganic chemicals react with HF rather than dissolving.[17]
Health effects [edit]
HF burns, not evident until a day later
Hydrogen fluoride is one of the toxic gases produced naturally by volcanic eruptions.[18]
Upon contact with moisture, including tissue, hydrogen fluoride immediately converts to hydrofluoric acid, which is highly corrosive and toxic. Exposure requires immediate medical attending.[xix] It can cause blindness by rapid destruction of the corneas. Breathing in hydrogen fluoride at loftier levels or in combination with peel contact can crusade death from an irregular heartbeat or from fluid buildup in the lungs.[19]
References [edit]
- ^ a b c d NIOSH Pocket Guide to Chemical Hazards. "#0334". National Institute for Occupational Prophylactic and Health (NIOSH).
- ^ Evans, D. A. "pKa's of Inorganic and Oxo-Acids" (PDF) . Retrieved June 19, 2020.
- ^ a b "Hydrogen fluoride". Immediately Dangerous to Life or Health Concentrations (IDLH). National Constitute for Occupational Safety and Wellness (NIOSH).
- ^ Johnson, M. W.; Sándor, E.; Arzi, Due east. (1975). "The Crystal Construction of Deuterium Fluoride". Acta Crystallographica. B31 (8): 1998–2003. doi:10.1107/S0567740875006711.
- ^ McLain, Sylvia Due east.; Benmore, C. J.; Siewenie, J. E.; Urquidi, J.; Turner, J. F. (2004). "On the Construction 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 Chemic Bail and the Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry . Cornell University Press. pp. 454–464. ISBN978-0-8014-0333-0.
- ^ Atkins, Peter; Jones, Loretta (2008). Chemical principles: The quest for insight. Westward. H. Freeman & Co. pp. 184–185. ISBN978-1-4292-0965-vi.
- ^ Emsley, John (1981). "The hidden forcefulness of hydrogen". New Scientist. 91 (1264): 291–292. Retrieved 25 December 2012.
- ^ Greenwood, N. Due north.; Earnshaw, A. (1998). Chemistry of the Elements (2nd ed.). Oxford: Butterworth Heinemann. pp. 812–816. ISBN0-7506-3365-4.
- ^ C. Due east. Housecroft and A. G. Sharpe Inorganic Chemistry, p. 221.
- ^ F. A. Cotton and 1000. Wilkinson Avant-garde Inorganic Chemical science, p. 111.
- ^ W. L. Jolly "Modern Inorganic Chemical science" (McGraw-Colina 1984), p. 203. ISBN 0-07-032768-8.
- ^ F. A. Cotton wool and M. Wilkinson, Advanced Inorganic Chemistry (5th ed.) John Wiley and Sons: New York, 1988. ISBN 0-471-84997-9. p. 109.
- ^ a b c d e J. Aigueperse, P. Mollard, D. Devilliers, 1000. Chemla, R. Faron, R. Romano, J. P. Cuer (2000). "Fluorine Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:x.1002/14356007.a11_307.
{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link) - ^ a b 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:ten.1002/14356007.a11_349.
{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link) - ^ M. Jaccaud, R. Faron, D. Devilliers, R. Romano (2005). "Fluorine". Ullmann'southward Encyclopedia of Industrial Chemical science. Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_293.
{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link). - ^ Greenwood and Earnshaw, "Chemistry of the Elements", pp. 816–819.
- ^ "Volcanoes". U.South. Environmental Protection Agency. Retrieved 20 December 2021.
- ^ a b Facts About Hydrogen Fluoride (Hydrofluoric Acid)
External links [edit]
- "ATSDR – Fluorides, Hydrogen Fluoride, and Fluorine". Retrieved September xxx, 2019
- CDC - NIOSH Pocket Guide to Chemical Hazards
- Toxics Use Reduction Institute - Hydrogen Fluoride Fact Sheet
Source: https://en.wikipedia.org/wiki/Hydrogen_fluoride
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