Hydrogen fluoride

Hydrogen fluoride

Names
Other names Fluorane
Identifiers
CAS Number	7664-39-3 Y
3D model (JSmol)	Interactive image
ChEBI	CHEBI:29228 Y
ChemSpider	14214 Y
ECHA InfoCard	100.028.759
KEGG	C16487 Y
PubChem CID	16211014
RTECS number	MW7875000
UNII	RGL5YE86CZ Y
UN number	1052
InChI InChI=1S/FH/h1H Y Key: KRHYYFGTRYWZRS-UHFFFAOYSA-N Y InChI=1/FH/h1H Key: KRHYYFGTRYWZRS-UHFFFAOYAC
SMILES F
Properties
Chemical formula	HF
Molar mass	7001200060000000000♠20.006 g·mol−1
Appearance	colourless gas or colourless liquid (below 19.5 °C)
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)
Solubility in water	completely miscible (liquid)
Vapor pressure	783 mmHg (20 °C)[1]
Acidity (pKa)	3.17[2][3]
Conjugate acid	Fluoronium
Conjugate base	Fluoride
Refractive index (nD)	1.00001
Structure
Molecular shape	Linear
Dipole moment	1.86 D
Thermochemistry
Std molar entropy (So298)	8.687 J/g K (gas)
Std enthalpy of formation (ΔfHo298)	−13.66 kJ/g (gas) −14.99 kJ/g (liquid)
Hazards
GHS pictograms
GHS signal word	Danger
GHS hazard statements	H300, H310, H314, H330
GHS precautionary statements	P260, P262, P264, P270, P271, P280, P284, P301+310, P301+330+331, P302+350, P303+361+353, P304+340, P305+351+338, P310, P320, P321, P322, P330, P361, P363, P403+233, P405, P501
NFPA 704	0 4 1
Lethal dose or concentration (LD, LC):
LC50 (median concentration)	1276 ppm (rat, 1 hr) 1774 ppm (monkey, 1 hr) 4327 ppm (guinea pig, 15 min)[4]
LCLo (lowest published)	313 ppm (rabbit, 7 hr)[4]
US health exposure limits (NIOSH):
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
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).
Y verify (what is YN ?)
Infobox references

Hydrogen fluoride is a chemical compound with the chemical formula HF. This colorless gas or liquid is the principal industrial source of fluorine, often as an aqueous solution called hydrofluoric acid. It is an important feedstock in the preparation of many important compounds including pharmaceuticals and polymers (e.g. Teflon). HF is widely used in the petrochemical industry as a component of superacids. Hydrogen fluoride boils near room temperature, much higher than other hydrogen halides.

Hydrogen fluoride is a highly dangerous gas, forming corrosive and penetrating hydrofluoric acid upon contact with moisture. The gas can also cause blindness by rapid destruction of the corneas.

French chemist Edmond Frémy (1814–1894) is credited with discovering anhydrous hydrogen fluoride while trying to isolate fluorine. Although Carl Wilhelm Scheele prepared hydrofluoric acid in large quantities in 1771, this acid was known in the glass industry before then.

Contents

• 
  1 Structure
  
  
  
  • 
    1.1 Comparison with other hydrogen halides
    
  
  
  
  


• 
  2 Acidity
  


• 
  3 Solvent
  


• 
  4 Production and uses
  


• 
  5 Health effects
  


• 
  6 References
  


• 
  7 External links
  

Structure

Although a diatomic molecule, HF forms relatively strong intermolecular hydrogen bonds. Solid HF consists of zigzag 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.^[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]

Comparison with other hydrogen halides

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).^[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.

Hydrogen fluoride is miscible with water (dissolves in any proportion), whereas the other hydrogen halides have large solubility gaps with water. Hydrogen fluoride and water also form several compounds in the solid state, most notably a 1:1 compound that does not melt until −40 °C (−40 °F), which is 44 °C (79 °F) above the melting point of pure HF.^[10]

HF and H2O similarities
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.

Acidity

Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in dilute aqueous solution.^[11] This is in part a result of the strength of the hydrogen–fluorine bond, but also of other factors such as the tendency of HF, H
₂O, and F⁻
anions to form clusters.^[12] At high concentrations, HF molecules undergo homoassociation to form polyatomic ions (such as bifluoride, HF⁻
₂) and protons, thus greatly increasing the acidity.^[13] This leads to protonation of very strong acids like hydrochloric, sulfuric, or nitric when using concentrated hydrofluoric acid solutions.^[14] Although hydrofluoric acid is regarded as a weak acid, it is very corrosive, even attacking glass when hydrated.^[13]

The acidity of hydrofluoric acid solutions varies with concentration owing to hydrogen-bond interactions of the fluoride ion. Dilute solutions are weakly acidic with an acid ionization constant K_a = 6.6×10⁻⁴ (or pK_a = 3.18),^[15] in contrast to corresponding solutions of the other hydrogen halides, which are strong acids (pK_a < 0). Concentrated solutions of hydrogen fluoride are much more strongly acidic than implied by this value, as shown by measurements of the Hammett acidity function H₀^[16](or "effective pH"). The H₀ for 100% HF is estimated to be between −10.2 and −11, comparable to the value −12 for sulfuric acid.^[17][18]

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.^[19] However, Paul Giguère and Sylvia Turrell^[20][21] have shown by infrared spectroscopy that the predominant solute species is the hydrogen-bonded ion pair [H
3O+
·F⁻], which suggests that the ionization can be described as a pair of successive equilibria:

H 2O + HF ↽−⇀{displaystyle {ce {<=>>}}} [H 3O+ ·F−]

(1)

[H 3O+ ·F−] ↽−⇀{displaystyle {ce {<<=>}}} H 3O+ + F−

(2)

The first equilibrium lies well to the right (K ≫ 1) and the second to the left (K ≪ 1), 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.^[22]

In concentrated solution, the additional HF causes the ion pair to dissociate with formation of the hydrogen-bonded hydrogen difluoride ion.^[20][22]

[H
3O+
⋅F⁻] + HF ⇌ H
3O+
+ HF⁻
₂

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 HF⁻
₂. This interaction between the acid and its own conjugate base is an example of homoassociation (homoconjugation). At the limit of 100% liquid HF, there is self-ionization^[23][24]

3 HF ⇌ H₂F⁺ + HF⁻
₂

which 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.^[17][18]

Solvent

Dry hydrogen fluoride readily dissolves low-valent metal fluorides, as well as several molecular fluorides. Many proteins and carbohydrates can be dissolved in dry HF and recovered from it. In contrast, most non-fluoride inorganic chemicals react with HF rather than dissolving.^[25]

Production and uses

Hydrogen fluoride is produced 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 petrol (gasoline) called "alkylate" is generated in alkylation units that combine C₃ and C₄ olefins and iso-butane to generate petrol (gasoline).^[26]

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.^[27]

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.^[28]

Acyl chlorides or acid anhydrides react with hydrogen fluoride to give acyl fluorides.^[29]

HF is often used in palynology to remove silicate minerals, for extraction of dinoflagellate cysts, acritarchs and chitinozoans.

1,1-Difluoroethane is produced by the mercury-catalyzed addition of hydrogen fluoride to acetylene:^[30]

HC≡CH + 2 HF → CH₃CHF₂

The intermediate in this process is vinyl fluoride or fluoroethylene, the monomeric precursor to polyvinyl fluoride.

Health effects

Main article: Hydrofluoric acid

HF burns, not evident until a day after

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.^[31] Breathing in hydrogen fluoride at high levels or in combination with skin contact can cause death from an irregular heartbeat or from fluid buildup in the lungs.^[31]

References

1. 
   ^ ^abcd NIOSH Pocket Guide to Chemical Hazards. "#0334". National Institute for Occupational Safety and Health (NIOSH)..mw-parser-output cite.citation{font-style:inherit}.mw-parser-output .citation q{quotes:"""""""'""'"}.mw-parser-output .citation .cs1-lock-free a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/6/65/Lock-green.svg/9px-Lock-green.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .citation .cs1-lock-limited a,.mw-parser-output .citation .cs1-lock-registration a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/d/d6/Lock-gray-alt-2.svg/9px-Lock-gray-alt-2.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .citation .cs1-lock-subscription a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/a/aa/Lock-red-alt-2.svg/9px-Lock-red-alt-2.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .cs1-subscription,.mw-parser-output .cs1-registration{color:#555}.mw-parser-output .cs1-subscription span,.mw-parser-output .cs1-registration span{border-bottom:1px dotted;cursor:help}.mw-parser-output .cs1-ws-icon a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/4/4c/Wikisource-logo.svg/12px-Wikisource-logo.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output code.cs1-code{color:inherit;background:inherit;border:inherit;padding:inherit}.mw-parser-output .cs1-hidden-error{display:none;font-size:100%}.mw-parser-output .cs1-visible-error{font-size:100%}.mw-parser-output .cs1-maint{display:none;color:#33aa33;margin-left:0.3em}.mw-parser-output .cs1-subscription,.mw-parser-output .cs1-registration,.mw-parser-output .cs1-format{font-size:95%}.mw-parser-output .cs1-kern-left,.mw-parser-output .cs1-kern-wl-left{padding-left:0.2em}.mw-parser-output .cs1-kern-right,.mw-parser-output .cs1-kern-wl-right{padding-right:0.2em}
   


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. 
   ^ ^ab "Hydrogen fluoride". Immediately Dangerous to Life and 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, 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.
   


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 978-1-4292-0965-6.
   


9. 
   ^ Emsley, John (1981). "The hidden strength of hydrogen". New Scientist. 91 (1264): 291–292. 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. 
    ^ Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic Chemistry. San Diego: Academic Press. p. 425. ISBN 978-0-12-352651-9.
    


12. 
    ^ Clark, Jim (2002). "The acidity of the hydrogen halides". Retrieved 4 September 2011.
    


13. 
    ^ ^ab Chambers, C.; Holliday, A. K. (1975). Modern inorganic chemistry (An intermediate text) (PDF). The Butterworth Group. pp. 328–329. Archived from the original (PDF) on 2013-03-23.
    


14. 
    ^ Hannan, Henry J. (2010). Course in chemistry for IIT-JEE 2011. Tata McGraw Hill Education Private Limited. pp. 15–22. ISBN 9780070703360.
    


15. 
    ^ 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.
    


16. 
    ^ Hyman H. H., Kilpatrick M., Katz J. J. (1957). "The Hammett Acidity Function H₀ for Hydrofluoric Acid Solutions". Journal of the American Chemical Society. 79 (14): 3668–3671. doi:10.1021/ja01571a016. ISSN 0002-7863.CS1 maint: Uses authors parameter (link)
    


17. 
    ^ ^ab W. L. Jolly “Modern Inorganic Chemistry” (McGraw-Hill 1984), p. 203.
    ISBN 0-07-032768-8.
    


18. 
    ^ ^ab F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry (5th ed.) John Wiley and Sons: New York, 1988.
    ISBN 0-471-84997-9. p. 109.
    


19. 
    ^ C. E. Housecroft and A. G. Sharpe "Inorganic Chemistry" (Pearson Prentice Hall, 2nd ed. 2005), p. 170.
    


20. 
    ^ ^ab Giguère, Paul A.; Turrell, Sylvia (1980). "The nature of hydrofluoric acid. A spectroscopic study of the proton-transfer complex H
    3O+
    ...F⁻". J. Am. Chem. Soc. 102 (17): 5473. doi:10.1021/ja00537a008.
    


21. 
    ^ Radu Iftimie; Vibin Thomas; Sylvain Plessis; Patrick Marchand; Patrick Ayotte (2008). "Spectral Signatures and Molecular Origin of Acid Dissociation Intermediates". J. Am. Chem. Soc. 130 (18): 5901–7. doi:10.1021/ja077846o. PMID 18386892.
    


22. 
    ^ ^ab F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, p. 104.
    


23. 
    ^ C. E. Housecroft and A. G. Sharpe Inorganic Chemistry, p. 221.
    


24. 
    ^ F. A. Cotton and G. Wilkinson Advanced Inorganic Chemistry, p. 111.
    


25. 
    ^ Greenwood and Earnshaw, "Chemistry of the Elements", pp. 816–819.
    


26. 
    ^ Aigueperse J, Mollard P, Devilliers D, Chemla M, Faron R, Romano R, Cuer JP (2000). "Fluorine Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a11_307. ISBN 3527306730.
    


27. 
    ^ G. Siegemund, W. Schwertfeger, A. Feiring, B. Smart, F. Behr, H. Vogel, B. McKusick "Fluorine Compounds, Organic" in "Ullmann’s Encyclopedia of Industrial Chemistry" 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a11_349
    


28. 
    ^ 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.
    


29. 
    ^ Olah G, Kuhn S (1961). "Preparation of Acyl Fluorides with Anhydrous Hydrogen Fluoride. The General Use of the Method of Colson and Fredenhagen". J. Org. Chem. 26: 237–238. doi:10.1021/jo01060a600.
    


30. 
    ^ Siegemund, Günter; Schwertfeger, Werner; Feiring, Andrew; Smart, Bruce; Behr, Fred; Vogel, Herward; McKusick, Blaine (2010). "Fluorine Compounds, Organic". In Bohnet, Matthias; Bellussi, Giuseppe; Bus, James; et al. Ullmann's Encyclopedia of Industrial Chemistry. John Wiley & Sons. doi:10.1002/14356007.a11_349.
    


31. 
    ^ ^ab Facts About Hydrogen Fluoride (Hydrofluoric Acid)
    

External links

Wikimedia Commons has media related to Hydrogen fluoride.

• 
  "ATSDR – MMG: Hydrogen Fluoride". Retrieved May 14, 2006


• CDC - NIOSH Pocket Guide to Chemical Hazards

v t e Hydrogen compounds

H3AsO3 H3AsO4 HArF HAt HSO3F HBF4 HBr HBrO HBrO2 HBrO3 HBrO4 HCl HClO HClO2 HClO3 HClO4 HCN HCNO H2CrO4/H2Cr2O7 H2CO3 H2CS3 HF HFΟ HI HIO HIO2 HIO3 HIO4 HMnO4 H2MoO4 HNC NaHCO3 HNCO HNO HNO2 HNO3 H2N2O2 HNO5S H3NSO3 H2O H2O2 H2O3 H3PO2 H3PO3 H3PO4 H4P2O7 H5P3O10 H2PtCl6 H2S H2S2 H2Se H2SeO3 H2SeO4 H4SiO4 H2SiF6 HSCN HNSC H2SO3 H2SO4 H2SO5 H2S2O3 H2S2O6 H2S2O7 H2S2O8 CF3SO3H H2Te H2TeO3 H6TeO6 H4TiO4 H2Po HCo(CO)4

v t e Molecules detected in outer space
Molecules	Diatomic Aluminium monochloride Aluminium monofluoride Aluminium monoxide Argonium Carbon monophosphide Carbon monosulfide Carbon monoxide Carborundum Cyanogen radical Diatomic carbon Fluoromethylidynium Hydrogen chloride Hydrogen fluoride Hydrogen (molecular) Hydroxyl radical Iron(II) oxide Magnesium monohydride cation Methylidyne radical Nitric oxide Nitrogen (molecular) Nitrogen monohydride Nitrogen sulfide Oxygen (molecular) Phosphorus monoxide Phosphorus mononitride Potassium chloride Silicon carbide Silicon mononitride Silicon monoxide Silicon monosulfide Sodium chloride Sodium iodide Sulfur monohydride Sulfur monoxide Titanium oxide Triatomic Aluminium hydroxide Aluminium isocyanide Amino radical Carbon dioxide Carbonyl sulfide CCP radical Chloronium Diazenylium Dicarbon monoxide Disilicon carbide Ethynyl radical Formyl radical Hydrogen cyanide (HCN) Hydrogen isocyanide (HNC) Hydrogen sulfide Hydroperoxyl Iron cyanide Isoformyl Magnesium cyanide Magnesium isocyanide Methylene radical N2H+ Nitrous oxide Nitroxyl Ozone Phosphaethyne Potassium cyanide Protonated molecular hydrogen Sodium cyanide Sodium hydroxide Silicon carbonitride c-Silicon dicarbide SiNC Sulfur dioxide Thioformyl Thioxoethenylidene Titanium dioxide Tricarbon Water Four atoms Acetylene Ammonia Cyanic acid Cyanoethynyl Cyclopropynylidyne Formaldehyde Fulminic acid HCCN Hydrogen peroxide Hydromagnesium isocyanide Isocyanic acid Isothiocyanic acid Ketenyl Methylene amidogen Methyl radical Propynylidyne Protonated carbon dioxide Protonated hydrogen cyanide Silicon tricarbide Thioformaldehyde Tricarbon monoxide Tricarbon sulfide Thiocyanic acid Five atoms Ammonium ion Butadiynyl Carbodiimide Cyanamide Cyanoacetylene Cyanoformaldehyde Cyanomethyl Cyclopropenylidene Formic acid Isocyanoacetylene Ketene Methane Methoxy radical Methylenimine Propadienylidene Protonated formaldehyde Protonated formaldehyde Silane Silicon-carbide cluster Six atoms Acetonitrile Cyanobutadiynyl radical E-Cyanomethanimine Cyclopropenone Diacetylene Ethylene Formamide HC4N Ketenimine Methanethiol Methanol Methyl isocyanide Pentynylidyne Propynal Protonated cyanoacetylene Seven atoms Acetaldehyde Acrylonitrile Vinyl cyanide Cyanodiacetylene Ethylene oxide Glycolonitrile Hexatriynyl radical Methylacetylene Methylamine Methyl isocyanate Vinyl alcohol Eight atoms Acetic acid Aminoacetonitrile Cyanoallene Ethanimine Glycolaldehyde Heptatrienyl radical Hexapentaenylidene Methylcyanoacetylene Methyl formate Propenal Nine atoms Acetamide Cyanohexatriyne Cyanotriacetylene Dimethyl ether Ethanol Methyldiacetylene Octatetraynyl radical Propene Propionitrile Ten atoms or more Acetone Benzene Benzonitrile Buckminsterfullerene (C60 fullerene, buckyball) C70 fullerene Cyanodecapentayne Cyanopentaacetylene Cyanotetra-acetylene Ethylene glycol Ethyl formate Methyl acetate Methyl-cyano-diacetylene Methyltriacetylene Propanal n-Propyl cyanide Pyrimidine
Deuterated molecules	Ammonia Ammonium ion Formaldehyde Formyl radical Heavy water Hydrogen cyanide Hydrogen deuteride Hydrogen isocyanide Methylacetylene N2D+ Trihydrogen cation
Unconfirmed	Anthracene Dihydroxyacetone Ethyl methyl ether Glycine Graphene H2NCO+ Linear C5 Naphthalene cation Phosphine Pyrene Silylidine
Related	Abiogenesis Astrobiology Astrochemistry Atomic and molecular astrophysics Chemical formula Circumstellar envelope Cosmic dust Cosmic ray Cosmochemistry Diffuse interstellar band Earliest known life forms Extraterrestrial life Extraterrestrial liquid water Forbidden mechanism Helium hydride ion Homochirality Intergalactic dust Interplanetary medium Interstellar medium Photodissociation region Iron–sulfur world theory Kerogen Molecules in stars Nexus for Exoplanet System Science Organic compound Outer space PAH world hypothesis Panspermia Polycyclic aromatic hydrocarbon (PAH) RNA world hypothesis Spectroscopy Tholin
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v t e Fluorine compounds
HF He LiF BeF2 BF3 B2F4 CF4 NF3 OF2 F− Ne NaF MgF2 AlF3 SiF4 PF3 PF5 SF4 SF6 ClF3 ClF5 Ar KF CaF2 ScF3 TiF3 TiF4 VF3 VF4 VF5 CrF2 CrF3 CrF4 CrF5 MnF2 MnF3 FeF2 FeF3 CoF2 CoF3 NiF2 CuF CuF2 ZnF2 GaF3 GeF4 AsF3 AsF5 SeF4 SeF6 BrF3 BrF5 KrF2 RbF SrF2 YF3 ZrF4 NbF4 NbF5 MoF5 MoF6 TcF6 RuF3 RuF4 RuF5 RuF6 RhF6 PbF2 AgF AgF2 Ag2F CdF2 InF3 SnF2 SnF4 SbF3 SbF5 TeF4 IF5 IF7 XeF2 XeF4 XeF6 CsF BaF2 * HfF4 TaF5 WF6 ReF4 ReF5 OsF4 OsF5 IrF3 PtF4 AuF3 AuF5 HgF2 Hg2F2 TlF TlF3 PbF2 PbF4 BiF3 BiF5 Po At Rn Fr Ra ** Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og ↓ * LaF3 CeF3 Pr NdF3 Pm Sm Eu Gd Tb DyF3 Ho Er Tm Yb Lu ** AcF3 ThF4 Pa UF4 UF6 UO2F2 Np Pu AmF3 AmF4 Cm Bk Cf Es Fm Md No Lr
Other	AgBF4 AgPF6 Cs2AlF5 K3AlF6 Na3AlF6 KAsF6 LiAsF6 NaAsF6 HBF4 KBF4 LiBF4 NaBF4 RbBF4 Ba(BF4)2 Ni(BF4)2 Pb(BF4)2 Sn(BF4)2 BaClF BaSiF6 BaGeF6 BrOF3 BrO2F CBrF3 CBr2F2 CBr3F CClF3 CCl2F2 CCl3F CFN CF2O CF3I CHF3 CH2F2 CH3F C2Cl3F3 C2H3F C6H5F C7H5F3 C15F33N ClFO2 CrFO4 CrF2O2 CsBF4 NH4F FNO FNO2 FNO3 KHF2 NaHF2 ThOF2 NH5F2 (NH4)2SiF6 F2OS F3OP PSF3 HPF6 HSbF6 NH4PF6 KPF6 KSbF6 LiPF6 NaPF6 NaSbF6 Na2SiF6 Na2TiF6 Na2ZrF6 TlPF6 IOF3 K2NbF7 K2TaF7 IO3F
Chemical formulas

v t e Binary compounds of hydrogen
Alkali metal hydrides	LiH NaH KH RbH CsH	Lithium hydride, LiH ionic metal hydride Beryllium hydride Left (gas phase): BeH2 covalent metal hydride Right: (BeH2)n (solid phase) polymeric metal hydride Borane and diborane Left: BH3 (special conditions), covalent metalloid hydride Right: B2H6 (standard conditions), dimeric metalloid hydride Methane, CH4 covalent nonmetal hydride Ammonia, NH3 covalent nonmetal hydride Water, H2O covalent nonmetal hydride Hydrogen fluoride, HF covalent nonmetal hydride
Alkaline earth hydrides	Monohydrides BeH MgH CaH SrH BaH Dihydrides BeH2 MgH2 CaH2 SrH2 BaH2
Group 13 hydrides	Boranes BH3 B2H6 B2H2 B2H4 B4H10 B5H9 B5H11 B6H10 B6H12 B10H14 B18H22 Alanes AlH3 Al2H6 Gallanes GaH3 Ga2H6 Indiganes InH3 In2H6 Thallanes TlH3 Tl2H6
Group 14 hydrides	Linear alkanes CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18 C9H20 C10H22 more... Linear alkenes C2H4 C3H6 C4H8 C5H10 C6H12 C7H14 C8H16 C9H18 C10H20 more... Linear alkynes C2H2 C3H4 C4H6 C5H8 C6H10 C7H12 C8H14 C9H16 C10H18 more... Silanes SiH4 Si2H6 Si3H8 Si4H10 Si5H12 Si6H14 Si7H16 Si8H18 Si9H20 Si10H22 more... Silenes Si2H4 Silynes Si2H2 SiH Germanes GeH4 Ge2H6 Ge3H8 Ge4H10 Ge5H12 Stannanes SnH4 Sn2H6 Plumbanes PbH4 CH CH2 CH3 C2H Cycloalkanes Cycloalkenes Annulenes Many more
Pnictogen hydrides	Azanes NH3 N2H4 N3H5 N4H6 N5H7 N6H8 N7H9 N8H10 N9H11 N10H12 more... Azenes N2H2 N3H3 N4H4 Phosphanes PH3 P2H4 P3H5 P4H6 P5H7 P6H8 P7H9 P8H10 P9H11 P10H12 more... Phosphenes P2H2 P3H3 P4H4 Arsanes AsH3 As2H4 Stibanes SbH3 Bismuthanes BiH3 HN3 NH radical
Hydrogen chalcogenides	Polyoxidanes H2O H2O2 H2O3 H2O4 H2O5 H2O6 H2O7 H2O8 H2O9 H2O10 more... Polysulfanes H2S H2S2 H2S3 H2S4 H2S5 H2S6 H2S7 H2S8 H2S9 H2S10 more... Selanes H2Se H2Se2 Tellanes H2Te H2Te2 Polanes PoH2 HO HO2 HO3 H2O+–O– H2S=S (HS)2S+–S– HS HDO D2O T2O
Hydrogen halides	HF HCl HBr HI HAt
Transition metal hydrides	ScH2 YH2 YH3 TiH2 ZrH2 HfH2 VH VH2 NbH NbH2 TaH CrH CrH2 CrHx NiH PdHx (x < 1) FeH FeH2 FeH5 CuH ZnH2 CdH2 HgH2
Lanthanide hydrides	LaH2 LaH3 LaH10 CeH2 CeH3 PrH2 PrH3 NdH2 NdH3 SmH2 SmH3 EuH2 GdH2 GdH3 TbH2 TbH3 DyH2 DyH3 HoH2 HoH3 ErH2 ErH3 TmH2 TmH3 YbH2 YbH2.5 LuH2 LuH3
Actinide hydrides	AcH2 ThH2 Th4H15 PaH3 UH3 NpH2 NpH3 PuH2 PuH3 AmH2 AmH3 CmH2

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