فلوريد الهيدروجين

(تم التحويل من Hydrogen fluoride)
Hydrogen fluoride
المُعرِّفات
رقم CAS
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.028.759 Edit this at Wikidata
KEGG
رقم RTECS
  • MW7875000
UNII
الخصائص
الصيغة الجزيئية HF
كتلة مولية 20.01 g mol-1
المظهر colourless gas or colourless liquid (below 19.5°C)
الكثافة 1.15 g/L, gas (25 °C)
0.99 g/mL, liquid (19.5 °C)
نقطة الانصهار
نقطة الغليان
قابلية الذوبان في الماء miscible
ضغط البخار 783 mmHg (20°C)[1]
الحموضة (pKa) 3.17[2][3]
معامل الانكسار (nD) 1.00001
البنية
الشكل الجزيئي Linear
Dipole moment 1.86 D
الكيمياء الحرارية
الإنتالپية المعيارية
للتشكل ΔfHo298 		−13.66 kJ/g (gas)
−14.99 kJ/g (liquid)
Standard molar
entropy So298 		8.687 J/g K (gas)
المخاطر
خطر رئيسي Highly toxic, corrosive, irritant
ن.م.ع. مخطط تصويري مصوّر التآكل في النظام المنسق عالمياً لتصنيف وعنونة الكيماويات (GHS) رمز الجمجمة والعظمتين في النظام المنسق عالمياً لتصنيف وعنونة الكيماويات (GHS)The exclamation-mark pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)رمز الخطر على الصحة في النظام المنسق عالمياً لتصنيف وعنونة الكيماويات (GHS)
ن.م.ع. كلمة الاشارة 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 (معيـَّن النار)
NFPA 704 four-colored diamondFlammability code 0: لن يشتعل. مثل الماءHealth code 4: التعرض لفترة قصيرة جداً قد يتسبب في الموت أو جروح بالغة باقية. مثل غاز VXReactivity code 1: مستقر في العادة، ولكن قد يصبح غير مستقر عند درجات الحرارة والضغط المرتفعين. مثال: الكالسيوم
0
4
1
نقطة الوميض none
الجرعة أو التركيز القاتل (LD, LC):
17 mg/kg (rat, oral)
1276 ppm (rat, 1 hr)
1774 ppm (monkey, 1 hr)
4327 ppm (guinea pig, 15 min)[4]
313 ppm (rabbit, 7 hr)[4]
حدود التعرض الصحية بالولايات المتحدة (NIOSH):
PEL (المسموح)
 TWA 3 ppm[1]
REL (الموصى به)
 TWA 3 ppm (2.5 mg/m3) C 6 ppm (5 mg/m3) [15 min][1]
IDLH (خطر عاجل)
 30 ppm[1]
مركبات ذا علاقة
Hydrogen chloride
Hydrogen bromide
Hydrogen iodide
Hydrogen astatide
Sodium fluoride
Potassium fluoride
Rubidium fluoride
Caesium fluoride
مركـّبات ذات علاقة
 Water
Ammonia
المخاطر
NFPA 704 (معيـَّن النار)
NFPA 704 four-colored diamondFlammability code 0: لن يشتعل. مثل الماءHealth code 4: التعرض لفترة قصيرة جداً قد يتسبب في الموت أو جروح بالغة باقية. مثل غاز VXReactivity code 1: مستقر في العادة، ولكن قد يصبح غير مستقر عند درجات الحرارة والضغط المرتفعين. مثال: الكالسيومSpecial hazards (white): no code
0
4
1
الجرعة أو التركيز القاتل (LD, LC):
1276 ppm (rat, 1 hr)
1774 ppm (monkey, 1 hr)
4327 ppm (guinea pig, 15 min)[4]
313 ppm (rabbit, 7 hr)[4]
حدود التعرض الصحية بالولايات المتحدة (NIOSH):
PEL (المسموح)
 TWA 3 ppm[1]
REL (الموصى به)
 TWA 3 ppm (2.5 mg/m3) C 6 ppm (5 mg/m3) [15-minute][1]
IDLH (خطر عاجل)
 30 ppm[1]
مركبات ذا علاقة
Hydrogen chloride
Hydrogen bromide
Hydrogen iodide
Sodium fluoride
Potassium fluoride
Rubidium fluoride
Caesium fluoride
مركـّبات ذات علاقة
 Water
Ammonia
ما لم يُذكر غير ذلك، البيانات المعطاة للمواد في حالاتهم العيارية (عند 25 °س [77 °ف]، 100 kPa).
YesY verify (what is YesYN ?)
مراجع الجدول

فلوريد الهيدروجين Hydrogen fluoride هو مركب كيميائي ذو الصيغة HF. وهو المصدر الصناعي الأساسي للفلور، ويكون غالبا بشكل حمض الهيدروفلوريك، فهو إذن الطليعة للعديد من المركبات المهمة مثل المركبات الصيدلانية ومبلمرات (مثل التيفلون). يستخدم فلوريد الهيدروجين استخدامًا واسعًا في صناعات البتروكيمياويات ومكون للعديد من الحموض الفائقة. يغلي فلوريد الهيدروجين في درجة حرارة أقل من درجة حرارة الغرفة في حين تتكاثف هاليدات الهيدروجين ‏(en)‏ الأخرى عند درجات حرارة أقل بكثير. وفلوريد الهيدروجين، خلافا لهاليدات الهيدروجين الأخرى، أخف من الهواء، وهو قادر على الاختراق خصوصا مما يجعله مؤذيا للرئتين. المحاليل المائية لفلوريد الهيدروجين، تسمى حمض الهيدروفلوريك، وهو مادة أكالة قوية.

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.

التاريخ

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 while trying to isolate fluorine.

البنية والتفاعلات

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 zigzag 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.[5] Liquid HF also consists of chains of HF molecules, but the chains are shorter, consisting on average of only five or six molecules.[6]

المقارنة مع هاليدات الهيدروجين الأخرى

Hydrogen fluoride does not boil until 20 °C (68 °F) in contrast to the heavier hydrogen halides, which boil between −85 و −35 °C (−121 و −31 °F).[7][8][9] This hydrogen bonding between HF molecules gives rise to high viscosity in the liquid phase and lower than expected pressure in the gas phase.

المحاليل المائية

المقالة الرئيسية: Hydrofluoric acid

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.[10]

HF and H2O similarities
graph showing trend-breaking water and HF boiling points: big jogs up versus a trend that is down with lower molecular weight for the other series members. graph showing humps of melting temperature, most prominent is at HF 50% mole fraction
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 [H
3O+
·F−].[11] However concentrated solutions are strong acids, because bifluoride anions are predominant, instead of ion pairs. In liquid anhydrous HF, self-ionization occurs:[12][13]

3 HF ⇌ H
2F+
 + HF−
2

which forms an extremely acidic liquid (H0 = −15.1).

التفاعلات مع أحماض لويس

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.[14][15]

التخليق

Hydrogen fluoride is typically produced by the reaction between sulfuric acid and pure grades of the mineral fluorite (calcium fluoride):[11][16]

CaF
2 + H
2SO
4 → 2 HF + CaSO
4

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:

H
2SiF
6 → 2 HF + SiF
4
SiF
4 + 2 H
2O → 4 HF + SiO
2

الاستخدام

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 synthetic cryolite for the electrolysis of aluminium.[16]

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):[16]

CHCl
3 + 2 HF → CHClF
2 + 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.[17]

1,1-Difluoroethane is produced by adding HF to acetylene using mercury as a catalyst.[17]

HC≡CH + 2 HF → CH
3CHF
2

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 aluminium produced. Other metal fluorides are produced using HF, including uranium tetrafluoride.[16]

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.[18]

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 isobutane.[16]

مذيب

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.[19]

التأثير على الصحة

left and right hands, two views, burned index fingers
HF burns, not evident until a day after
المقالة الرئيسية: Hydrofluoric acid and Hydrofluoric acid burn

Hydrogen fluoride is highly corrosive and a powerful contact poison. Exposure requires immediate medical attention.[20] 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).[20] Exposure of the intestinal system to HF solution is known to cause fulminant acute colitis requiring surgical intervention.[21]

الحموضة

تكون درجة الحموضة عالية ما لم يخفف بالماء.

المراجع

  1. ^ أ ب ت ث ج ح خ NIOSH Pocket Guide to Chemical Hazards 0334
  2. ^ "pKa's of Inorganic and Oxo-Acids" (PDF). Harvard. Retrieved 9 September 2013.
  3. ^ Bruckenstein, S.; Kolthoff, I.M., in Kolthoff, I.M.; Elving, P.J. Treatise on Analytical Chemistry, Vol. 1, pt. 1; Wiley, NY, 1959, pp. 432-433.
  4. ^ أ ب ت ث "Hydrogen fluoride". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  5. ^ Johnson, M. W.; Sándor, E.; Arzi, E. (1975). "The Crystal Structure of Deuterium Fluoride". Acta Crystallographica. B31 (8): 1998–2003. doi:10.1107/S0567740875006711.
  6. ^ 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): 1952–1955. doi:10.1002/anie.200353289. PMID 15065271.
  7. ^ 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.
  8. ^ Atkins, Peter; Jones, Loretta (2008). Chemical principles: The quest for insight. W. H. Freeman & Co. pp. 184–185. ISBN 978-1097774678.
  9. ^ Emsley, John (1981). "The hidden strength of hydrogen". New Scientist. 91 (1264): 291–292. Archived from the original on 22 July 2023. Retrieved 25 December 2012.
  10. ^ Greenwood, N. N.; Earnshaw, A. (1998). Chemistry of the Elements (2nd ed.). Oxford: Butterworth Heinemann. pp. 812–816. ISBN 0-7506-3365-4.
  11. ^ أ ب Rennie, Richard, ed. (2020). Dictionary of chemistry. Oxford quick reference (8th ed.). Oxford, United Kingdom ; New York, NY: Oxford University Press. ISBN 978-0-19-884122-7.
  12. ^ Housecroft, C. E.; Sharpe, A. G. Inorganic Chemistry. p. 221.[edition needed][ISBN missing]
  13. ^ Cotton, F. A.; Wilkinson, G. Advanced Inorganic Chemistry. p. 111.[edition needed][ISBN missing]
  14. ^ Jolly, W. L. (1984). Modern Inorganic Chemistry. McGraw-Hill. p. 203. ISBN 0-07-032768-8..
  15. ^ Cotton, F. A.; Wilkinson, G. (1988). Advanced Inorganic Chemistry (5th ed.). New York, NY: John Wiley and Sons. p. 109. ISBN 0-471-84997-9.
  16. ^ أ ب ت ث ج Aigueperse, J.; Mollard, P.; Devilliers, D.; Chemla, M.; Faron, R.; Romano, R.; Cuer, J. P. (2000). "Fluorine Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_307. {{cite encyclopedia}}: Cite has empty unknown parameter: |authors= (help)
  17. ^ أ ب Siegemund, G.; Schwertfeger, W.; Feiring, A.; Smart, B.; Behr, F.; Vogel, H.; McKusick, B. (2005). "Fluorine Compounds, Organic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_349. {{cite encyclopedia}}: Cite has empty unknown parameter: |authors= (help)
  18. ^ Jaccaud, M.; Faron, R.; Devilliers, D.; Romano, R. (2005). "Fluorine". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_293. {{cite encyclopedia}}: Cite has empty unknown parameter: |authors= (help).
  19. ^ Greenwood; Earnshaw. Chemistry of the Elements. pp. 816–819.[edition needed][ISBN missing]
  20. ^ أ ب "Facts About Hydrogen Fluoride (Hydrofluoric Acid)". Emergency Preparedness and Response. U.S. Centers for Disease Control and Prevention.
  21. ^ Cappell, M. S.; Simon, T. (January 1993). "Fulminant acute colitis following a self-administered hydrofluoric acid enema". The American Journal of Gastroenterology. 88 (1): 122–126. ISSN 0002-9270. PMID 8420252.

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