View source for Hydrogen fluoride

{{distinguish|text = the element [[hafnium]], symbol Hf}}
{{chembox
| verifiedrevid  = 476992809
| Name           = Hydrogen fluoride
| ImageFileL1    = Hydrogen-fluoride-2D-dimensions.svg
| ImageFile      = Hydrogen fluoride.svg
| ImageFileR1    = Hydrogen-fluoride-3D-vdW.svg
| OtherNames     = Fluorane
| Section1       = {{Chembox Identifiers
|UNII_Ref = {{fdacite|correct|FDA}}
|UNII = RGL5YE86CZ
|KEGG_Ref = {{keggcite|correct|kegg}}
|KEGG = C16487
|InChI = 1/FH/h1H
|ChEBI_Ref = {{ebicite|correct|EBI}}
|ChEBI = 29228
|SMILES = F
|InChIKey = KRHYYFGTRYWZRS-UHFFFAOYAC
|StdInChI_Ref = {{stdinchicite|correct|chemspider}}
|StdInChI = 1S/FH/h1H
|StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
|StdInChIKey = KRHYYFGTRYWZRS-UHFFFAOYSA-N
|CASNo = 7664-39-3
|CASNo_Ref = {{cascite|correct|CAS}}
|ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
|ChemSpiderID = 14214
|UNNumber = 1052
|PubChem = 16211014
|RTECS = MW7875000
}}
| Section2       = {{Chembox Properties
|H=1 | F=1
|Formula = HF
|Appearance = colourless gas or colourless liquid (below 19.5 °C)
|Odor = unpleasant
|Density = 1.15 g/L, gas (25 °C)<br />0.99 g/mL, liquid (19.5 °C)<br />1.663 g/mL, solid (−125 °C)
|Solubility = miscible (liquid)
|MeltingPtC = −83.6
|BoilingPtC = 19.5
|RefractIndex = 1.00001
|pKa = 3.17 (in water), 
15 (in DMSO) <ref>{{Cite web|last=Evans|first=D. A.|title=pKa's of Inorganic and Oxo-Acids|url=http://ccc.chem.pitt.edu/wipf/MechOMs/evans_pKa_table.pdf|access-date=June 19, 2020}}</ref>
|ConjugateAcid = [[Fluoronium]]
|ConjugateBase = [[Fluoride]]
|VaporPressure = 783 mmHg (20 °C)<ref name="PGCH" />
}}
| Section3       = {{Chembox Structure
|MolShape = [[Linear molecular geometry|Linear]]
|Dipole = 1.86 [[Debye|D]]
}}
| Section4       = {{Chembox Thermochemistry
|DeltaHf = −13.66 kJ/g (gas) <br /> −14.99 kJ/g (liquid)
|Entropy = 8.687 J/g K (gas)
}}
| Section5       = {{Chembox Hazards
|MainHazards = Highly toxic, corrosive, irritant
|GHSPictograms = {{GHS corrosion}} {{GHS skull and crossbones}}{{GHS07}}{{GHS health hazard}}
|GHSSignalWord = Danger
|HPhrases = {{H-phrases|300+310+330|314}}
|PPhrases = {{P-phrases|260|262|264|270|271|280|284|301+310|301+330+331|302+350|303+361+353|304+340|305+351+338|310|320|321|322|330|361|363|403+233|405|501}}
|NFPA-H = 4
|NFPA-F = 0
|NFPA-R = 1
|NFPA-S = POI
|FlashPt = none
|IDLH = 30 ppm<ref name="PGCH">{{PGCH|0334}}</ref>
|REL = TWA 3 ppm (2.5 mg/m<sup>3</sup>) C 6 ppm (5 mg/m<sup>3</sup>) [15-minute]<ref name="PGCH" />
|PEL = TWA 3 ppm<ref name="PGCH" />
|LD50 = 17 ppm (rat, oral)
|LC50 = 1276 ppm (rat, 1 hr)<br/>1774 ppm (monkey, 1 hr)<br/>4327 ppm (guinea pig, 15 min)<ref name="IDLH">{{IDLH|7664393|Hydrogen fluoride}}</ref>
|LCLo = 313 ppm (rabbit, 7 hr)<ref
name="IDLH"/>
}}
| Section6       = {{Chembox Related
|OtherAnions = [[Hydrogen chloride]]<br />[[Hydrogen bromide]]<br />[[Hydrogen iodide]]<br />[[Hydrogen astatide]]
|OtherCations = [[Sodium fluoride]]<br />[[Potassium fluoride]]<br />[[Rubidium fluoride]]<br />[[Caesium fluoride]]
|OtherCompounds = [[Properties of water|Water]]<br /> [[Ammonia]]
}}
}}

'''Hydrogen fluoride''' (fluorane) is an [[Inorganic chemistry|inorganic compound]] with [[chemical formula]] {{Chem2|HF|auto=yes}}. It is a very poisonous, colorless gas or liquid that dissolves in water to yield [[hydrofluoric acid]]. It is the principal industrial source of [[fluorine]], often in the form of hydrofluoric acid, and is an important [[feedstock]] in the preparation of many important compounds including pharmaceuticals and [[polymer]]s such as [[polytetrafluoroethylene]] (PTFE). HF is also widely used in the [[petrochemical industry]] as a component of [[superacid]]s. Due to strong and extensive [[hydrogen bond]]ing, it boils near room temperature, a much higher temperature than other [[hydrogen halide]]s.

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 [[cornea]]s.

==History==
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 (HF) while trying to isolate [[fluorine]].

==Structure and reactions==
[[File:Hydrogen-fluoride-solid-2D-dimensions.png|thumb|left|295px|The structure of chains of HF in crystalline hydrogen fluoride.]]{{clear|left}}
HF is diatomic in the gas-phase. As a liquid, HF forms relatively strong [[hydrogen bond]]s, 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&nbsp;pm length, are linked to neighboring molecules by intermolecular H–F distances of 155 pm.<ref>{{cite journal |author1=Johnson, M. W. |author2=Sándor, E. |author3=Arzi, E. | title = The Crystal Structure of Deuterium Fluoride | journal = [[Acta Crystallographica]] | year = 1975 | volume = B31 | pages = 1998–2003 | doi = 10.1107/S0567740875006711 | issue = 8 }}</ref>  Liquid HF also consists of chains of HF molecules, but the chains are shorter, consisting on average of only five or six molecules.<ref>{{cite journal | title = On the Structure of Liquid Hydrogen Fluoride | journal = [[Angewandte Chemie International Edition]] | year = 2004 | volume = 43 | pages = 1952–55 | doi = 10.1002/anie.200353289 | author = McLain, Sylvia E. | pmid = 15065271 | last2 = Benmore | first2 = C. J. | last3 = Siewenie | first3 = J. E. | last4 = Urquidi | first4 = J. | last5 = Turner | first5 = J. F. | issue = 15 }}</ref>

===Comparison with other hydrogen halides===
Hydrogen fluoride does not boil until 20&nbsp;°C in contrast to the heavier hydrogen halides, which boil between −85&nbsp;°C (−120&nbsp;°F) and −35&nbsp;°C (−30&nbsp;°F).<ref name="Pauling HF hydrogen bonds">{{cite book|last=Pauling|first=Linus A.|title=The Nature of the Chemical Bond and the Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry|year=1960|publisher=Cornell University Press|isbn=978-0-8014-0333-0|pages=[https://archive.org/details/natureofchemical00paul/page/454 454]–464|url=https://archive.org/details/natureofchemical00paul|url-access=registration}}</ref><ref name="Atkins HF">{{cite book|last=Atkins|first=Peter|title=Chemical principles: The quest for insight|year=2008|publisher=W. H. Freeman & Co|isbn=978-1097774678|pages=184–185|url=https://books.google.com/books?id=4R6hb1OIMRUC&pg=PA184|author2=Jones, Loretta}}</ref><ref name="New Scientist HF">{{cite journal|last=Emsley|first=John|title=The hidden strength of hydrogen|journal=New Scientist|year=1981|volume=91|issue=1264|pages=291–292|url=https://books.google.com/books?id=ZbthaZCUXy4C&pg=PA292|access-date=25 December 2012|archive-date=22 July 2023|archive-url=https://web.archive.org/web/20230722094703/https://books.google.com/books?id=ZbthaZCUXy4C&pg=PA292|url-status=dead}}</ref> This hydrogen bonding between HF molecules gives rise to high [[viscosity]] in the liquid phase and lower than expected pressure in the gas phase.

===Aqueous solutions===
{{main|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<sup>.</sup>H<sub>2</sub>O with melting point −40&nbsp;°C (−40&nbsp;°F), which is 44&nbsp;°C (79&nbsp;°F) above the melting point of pure HF.<ref>{{cite book|last1=Greenwood|first1=N. N.|last2=Earnshaw|first2=A.|title=Chemistry of the Elements|date=1998|edition=2nd|publisher=Butterworth Heinemann|isbn=0-7506-3365-4|location=Oxford|pages=812–816}}</ref>

{| style="margin:1em auto 1em 1em; text-align:left; width:505px; float:center;" cellpadding="0" cellspacing="0"
|-
| colspan="2" style="text-align:center;"| '''HF and H<sub>2</sub>O similarities'''
|-
| style="width:285px; padding-right:5px;" | [[File:Boiling-points Chalcogen-Halogen.svg|285px|alt=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.]]
| style="width:200px; padding-left:5px;" | [[File:HF-H2O Phase-Diagram.svg|200px|alt=graph showing humps of melting temperature, most prominent is at HF 50% mole fraction]]
|- style="vertical-align:top; font-size:.9em;"
| style="width:285px; padding-right:5px; padding-left:25px; line-height:1.4em; "|Boiling points of the hydrogen halides (blue) and [[hydrogen chalcogenide]]s (red): HF and H<sub>2</sub>O break trends.
| style="width:200px; padding-left:25px; line-height:1.4em; "|Freezing point of HF/ H<sub>2</sub>O 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 pair]]s [{{H3O+}}·F<sup>−</sup>]. However concentrated solutions are strong acids, because [[bifluoride]] anions are predominant, instead of ion pairs. In liquid anhydrous HF, [[molecular autoionization|self-ionization]] occurs:<ref>C. E. Housecroft and A. G. Sharpe ''Inorganic Chemistry'', p. 221.</ref><ref>F. A. Cotton and G. Wilkinson ''Advanced Inorganic Chemistry'', p. 111.</ref>

:{{chem2|3 HF <-> H2F+ + HF2-}}

which forms an extremely acidic liquid ({{math|[[Hammett acidity function|''H''<sub>0</sub>]]}}{{math|1=&nbsp;=&nbsp;−15.1}}).

===Reactions with Lewis acids===
Like water, HF can act as a weak base, reacting with [[Lewis acid]]s to give [[superacid]]s. A [[Hammett acidity function]] (''H''<sub>0</sub>) of −21 is obtained with [[antimony pentafluoride]] (SbF<sub>5</sub>), forming [[fluoroantimonic acid]].<ref name="Jolly">W. L. Jolly "Modern Inorganic Chemistry" (McGraw-Hill 1984), p. 203. {{ISBN|0-07-032768-8}}.</ref><ref name="Cotton 109">[[F. Albert Cotton|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.</ref>

==Production==
Hydrogen fluoride is typically produced by the reaction between [[sulfuric acid]] and pure grades of the mineral [[fluorite]]:<ref name="AigueperseMollard2000" />

:{{chem2|CaF2 + H2SO4 -> 2 HF + CaSO4}}

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:

:{{chem2|H2SiF6 -> 2 HF + SiF4}}
:{{chem2|SiF4 + 2 H2O -> 4 HF + SiO2}}

==Use==
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 compound]]s and a precursor to [[cryolite]] for the electrolysis of aluminium.<ref name="AigueperseMollard2000" />

===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):<ref name="AigueperseMollard2000" />

:{{chem2|CHCl3 + 2 HF -> CHClF2 + 2 HCl}}

Pyrolysis of chlorodifluoromethane (at 550-750&nbsp;°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 [[Carbon–fluorine bond|C–F bond]]s. [[Perfluorinated carboxylic acid]]s and [[sulfonic acid]]s are produced in this way.<ref name="Ullmann">{{Ullmann|author=G. Siegemund, W. Schwertfeger, A. Feiring, B. Smart, F. Behr, H. Vogel, B. McKusick|title=Fluorine Compounds, Organic|year=2005|doi=10.1002/14356007.a11_349}}</ref>

[[1,1-Difluoroethane]] is produced by adding HF to [[acetylene]] using mercury as a catalyst.<ref name="Ullmann" />
:{{chem2|HC\tCH + 2 HF -> CH3CHF2}}
The intermediate in this process is [[vinyl fluoride]] or fluoroethylene, the [[monomer]]ic 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 Al produced.  Other metal fluorides are produced using HF, including [[uranium tetrafluoride]].<ref name="AigueperseMollard2000" />

HF is the precursor to elemental [[fluorine]], F<sub>2</sub>, 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 F<sub>2</sub> are produced annually.<ref>{{Ullmann|author=M. Jaccaud, R. Faron, D. Devilliers, R. Romano|title=Fluorine|year=2005|doi=10.1002/14356007.a11_293}}.</ref>

===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 C<sub>3</sub> and C<sub>4</sub> olefins and [[methylpropane|''iso''-butane]].<ref name="AigueperseMollard2000">{{Ullmann|author = J. Aigueperse, P. Mollard, D. Devilliers, M. Chemla, R. Faron, R. Romano, J. P. Cuer|year=2000|doi=10.1002/14356007.a11_307|title=Fluorine Compounds, Inorganic|isbn = 3527306730}}</ref>

===Solvent===
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.<ref>Greenwood and Earnshaw, "Chemistry of the Elements", pp. 816–819.</ref>

==Health effects==
[[File:HF burned hands.jpg|thumb|250px|alt=left and right hands, two views, burned index fingers|HF burns, not evident until a day after]]
{{Main|Hydrofluoric acid|Hydrofluoric acid burn}}

Hydrogen fluoride is highly corrosive and a powerful contact poison. Exposure requires immediate medical attention.<ref name="emergency.cdc.gov">[http://emergency.cdc.gov/agent/hydrofluoricacid/basics/facts.asp Facts About Hydrogen Fluoride (Hydrofluoric Acid)]</ref> It can cause blindness by rapid destruction of the [[cornea]]s. Breathing in hydrogen fluoride at high levels or in combination with skin contact can cause death from an [[Cardiac dysrhythmia|irregular heartbeat]] or from [[pulmonary edema]] (fluid buildup in the lungs).<ref name="emergency.cdc.gov" />

{{Clear}}

== References ==
{{Reflist|30em}}

==External links==
{{Commons category|Hydrogen fluoride}}
*[https://www.atsdr.cdc.gov/substances/toxsubstance.asp?toxid=38 Fluorides, Hydrogen Fluoride, and Fluorine] at [[ATSDR]]. Retrieved September 30, 2019
*[https://www.cdc.gov/niosh/npg/npgd0334.html CDC - NIOSH Pocket Guide to Chemical Hazards]
*[https://web.archive.org/web/20191001001621/https://www.turi.org/TURI_Publications/TURI_Chemical_Fact_Sheets/Hydrogen_Fluoride_Fact_Sheet Hydrogen Fluoride Fact Sheet] at [[Toxics Use Reduction Institute]]

{{Hydrogen compounds}}
{{Molecules detected in outer space}}
{{fluorine compounds}}
{{fluorides}}
{{Hydrides by group}}
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[[Category:Fluorides]]
[[Category:Hydrogen compounds]]
[[Category:Industrial gases]]
[[Category:Inorganic solvents]]
[[Category:Nonmetal halides]]
[[Category:Diatomic molecules]]