View source for Hydrofluoric acid

{{short description|Solution of hydrogen fluoride in water}}
{{chembox
| IUPACName      = Fluorane<ref>{{cite book|title=Nomenclature of Organic Chemistry: IUPAC Recommendations and Preferred Names 2013|editor-last=Favre |editor-first=Henri A. |editor2-first=Warren H.|editor2-last=Powell|date=2014|publisher=The [[Royal Society of Chemistry]]|place=Cambridge|isbn=9781849733069 | page=131}}</ref>
| Watchedfields  = changed
| verifiedrevid  = 464361624
| Name           = Hydrofluoric acid (HF + H<sub>2</sub>O {{Eqm}} F<sup>-</sup> + H<sub>3</sub>O<sup>+</sup>)
| ImageFile3     = Hydrogen fluoride.JPG
| ImageClass3 = bg-transparent
| ImageFileL1    = Hydrogen-fluoride-3D-vdW.png
| ImageClassL1 = bg-transparent
| ImageSize1     = 100px
| ImageFileR1    = Water-3D-vdW.png
| ImageClassR1 = bg-transparent
| ImageFileR2    = Oxonium-ion-3D-vdW.png
| ImageClassR2 = bg-transparent
| ImageFileL2    = Fluoride_ion2.svg
| ImageClassL2 = bg-transparent
| ImageSize2     = 100px
| OtherNames     = Fluorhydric acid<br />Hydronium fluoride
| SystematicName = 
| Section1       = {{Chembox Identifiers
| CASNo = 7664-39-3
| CASNo_Ref = {{cascite|correct|CAS}}
|    RTECS  =  MW7875000
|  PubChem  =  14917
|  ChemSpiderID_Ref  = {{chemspidercite|correct|chemspider}}|  ChemSpiderID  =  14214
|  UNII_Ref  = {{fdacite|correct|FDA}}|  UNII  =  RGL5YE86CZ
|  InChI  =  1/FH/h1H
|  ChEBI_Ref  = {{ebicite|correct|EBI}}|  ChEBI  =  29228
|  StdInChI_Ref  = {{stdinchicite|correct|chemspider}}|  StdInChI  =  1S/FH/h1H
|  StdInChIKey_Ref  = {{stdinchicite|correct|chemspider}}|  StdInChIKey  =  KRHYYFGTRYWZRS-UHFFFAOYSA-N
|  SMILES  =  F.O
| SMILES1 = [F-].[OH3+]
|  InChIKey  =  KRHYYFGTRYWZRS-UHFFFAOYAC
|  EINECS  =  231-634-8
}}
| Section2       = {{Chembox Properties
| Formula = HF (aq)
| Appearance = Colorless liquid
| Density = 1.15 g/mL (for 48% soln.)
| pKa = 3.17<ref>{{cite book |last=Harris |first=Daniel C. |year=2010 |title=Quantitative Chemical Analysis |location=New York |publisher=W. H. Freeman |edition=8th international |pages=AP14 |isbn=978-1429263092 }}</ref>
  }}
| Section3       = 
| Section4       = 
| Section5       = 
| Section6       = 
| Section7       = {{Chembox Hazards
| Hazards_ref = <ref>{{Cite web|url=https://pubchem.ncbi.nlm.nih.gov/compound/hydrofluoric_acid|title=Hydrofluoric Acid|website=PubChem|publisher=National Institute of Health|access-date=October 12, 2017}}</ref>
| NFPA-H = 4
| NFPA-F = 0
| NFPA-R = 0
| NFPA-S =acid
| GHSPictograms = {{GHS05|Corrosive}}{{GHS06|Acute Toxicity}}
| GHSSignalWord = DANGER
| HPhrases = {{H-phrases|280|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|410+403|501}}
  }}
}}

'''Hydrofluoric acid''' is a [[Solution (chemistry)|solution]] of [[hydrogen fluoride]] (HF) in [[water]]. Solutions of HF are colorless, [[acid]]ic and highly [[Corrosive substance|corrosive]]. A common concentration is 49% (48–52%) but there are also stronger solutions (e.g. 70%) and pure HF has a boiling point near room temperature. It is used to make most [[organofluorine compound]]s; examples include the commonly used pharmaceutical antidepressant medication [[fluoxetine]] (Prozac) and the  material [[polytetrafluoroethylene|PTFE]] (Teflon).  Elemental [[fluorine]] is produced from it. It is commonly used to [[Etching (microfabrication)|etch]] glass and silicon wafers.

==Uses==
===Production of organofluorine compounds===
The principal use of hydrofluoric acid is in [[organofluorine chemistry]]. Many [[organofluorine compounds]] are prepared using HF as the fluorine source, including [[Polytetrafluoroethylene|Teflon]], [[fluoropolymer]]s, [[fluorocarbon]]s, and [[refrigeration|refrigerants]] such as [[freon]].  Many pharmaceuticals contain fluorine.<ref name=Ullmann/>

===Production of inorganic fluorides===
Most high-volume inorganic fluoride compounds are prepared from hydrofluoric acid.  Foremost are Na<sub>3</sub>AlF<sub>6</sub>, [[cryolite]], and AlF<sub>3</sub>, [[aluminium trifluoride]].  A molten mixture of these solids serves as a high-temperature solvent for the production of metallic [[aluminium]].  Other inorganic fluorides prepared from hydrofluoric acid include [[sodium fluoride]] and [[uranium hexafluoride]].<ref name=Ullmann/>

===Etchant, cleaner===
[[File:Wet etching tanks at LAAS (6 inches) 0468.jpg|thumbnail|upright=1.4|Wet etching tanks]]
It is used in the [[semiconductor]] industry as a major component of [[Wright etch]] and [[buffered oxide etch]], which are used to clean [[silicon wafer]]s. In a similar manner it is also used to [[Glass etching|etch glass]] by treatment with [[silicon dioxide]] to form gaseous or water-soluble silicon fluorides. It can also be used to polish and frost glass.<ref name=":0" />

:SiO<sub>2</sub> + 4 HF → [[Silicon tetrafluoride|SiF<sub>4</sub>]](g) + 2 H<sub>2</sub>O
:SiO<sub>2</sub> + 6 HF → [[Hexafluorosilicic acid|H<sub>2</sub>SiF<sub>6</sub>]] + 2 H<sub>2</sub>O

A 5% to 9% hydrofluoric acid gel is also commonly used to etch all ceramic dental restorations to improve bonding.<ref>{{cite book | last=Craig | first=Robert | title=Craig's restorative dental materials | publisher=Mosby Elsevier | location=St. Louis, Mo | year=2006 | isbn=0-323-03606-6 | oclc=68207297}}<!-- Powers, John M. and Sakaguchi, Ronald L. (2006) ''Craig's Restorative Dental Materials'', 12th ed., Mosby, {{ISBN|0323036066}} p. 10.5.3--></ref> For similar reasons, dilute hydrofluoric acid is a component of household rust stain remover, in [[car wash]]es in "wheel cleaner" compounds, in ceramic and fabric rust inhibitors, and in water spot removers.<ref name=":0" /><ref>{{cite journal|last=Strachan |first=John |title=A deadly rinse: The dangers of hydrofluoric acid |journal=Professional Carwashing & Detailing |date=January 1999 |url=http://secure.gvmg.com/carwash/articleprint.asp?print=1&IndexID=4230101 |volume=23 |issue=1 |url-status=dead |archive-url=https://web.archive.org/web/20120425234944/http://secure.gvmg.com/carwash/articleprint.asp?print=1&IndexID=4230101 |archive-date=April 25, 2012 }}</ref> Because of its ability to dissolve iron oxides as well as silica-based contaminants, hydrofluoric acid is used in pre-commissioning boilers that produce high-pressure steam.  Hydrofluoric acid is also useful for dissolving rock samples (usually powdered) prior to analysis. In similar manner, this acid is used in [[acid maceration]]s to extract organic fossils from silicate rocks.  Fossiliferous rock may be immersed directly into the acid, or a [[cellulose nitrate]] film may be applied (dissolved in [[amyl acetate]]), which adheres to the organic component and allows the rock to be dissolved around it.<ref name=Edwards1982>{{Cite journal
| last = Edwards | first =  D.
| year = 1982
| title = Fragmentary non-vascular plant microfossils from the late Silurian of Wales
| journal = Botanical Journal of the Linnean Society
| volume = 84
| issue = 3
| pages = 223–256
| doi = 10.1111/j.1095-8339.1982.tb00536.x
}}</ref>

===Oil refining===
In a standard [[oil refinery]] process known as [[alkylation]], [[isobutane]] is alkylated with low-molecular-weight [[alkene]]s (primarily a mixture of [[propylene]] and [[butylene]]) in the presence of an acid [[catalyst]] derived from hydrofluoric acid. The catalyst protonates the alkenes (propylene, butylene) to produce reactive [[carbocation]]s, which alkylate isobutane. The reaction is carried out at mild temperatures (0 and 30&nbsp;°C) in a two-phase reaction.

==Production==

Hydrofluoric acid was first prepared in 1771, by [[Carl Wilhelm Scheele]].<ref>{{Greenwood&Earnshaw1st|page=921}}</ref> It is now mainly produced by treatment of the mineral [[fluorite]], CaF<sub>2</sub>, with concentrated [[sulfuric acid]] at approximately 265&nbsp;°C.

:CaF<sub>2</sub>  +  H<sub>2</sub>SO<sub>4</sub>  →  2 HF  +  CaSO<sub>4</sub>

The acid is also a by-product of the production of [[phosphoric acid]] from [[apatite]] and [[fluoroapatite]]. Digestion of the mineral with sulfuric acid at elevated temperatures releases a mixture of gases, including hydrogen fluoride, which may be recovered.<ref name=Ullmann>{{Ullmann|doi=10.1002/14356007.a11_307|title=Fluorine Compounds, Inorganic|year=2000|last1=Aigueperse|first1=Jean|last2=Mollard|first2=Paul|last3=Devilliers|first3=Didier|last4=Chemla|first4=Marius|last5=Faron|first5=Robert|last6=Romano|first6=René|last7=Cuer|first7=Jean Pierre|isbn=3527306730}}</ref>

Because of its high reactivity toward glass, hydrofluoric acid is stored in fluorinated [[plastic]] (often [[PTFE]]) containers.<ref name=Ullmann/><ref name=":0">{{Cite web|work=  Emergency Response Safety and Health Database |title=Hydrogen Fluoride/Hydrofluoric Acid: Systemic Agent |url = https://www.cdc.gov/niosh/ershdb/EmergencyResponseCard_29750030.html|publisher= NIOSH - CDC |access-date = 2015-12-04 |date= May 12, 2011 |url-status=live |archive-url=https://web.archive.org/web/20151207082035/https://www.cdc.gov/niosh/ershdb/EmergencyResponseCard_29750030.html |archive-date= Dec 7, 2015 }}</ref>

==Properties==

In dilute aqueous solution, hydrogen fluoride behaves as a weak acid.<ref name="Pearson/Prentice Hall">{{cite book |author1=Ralph H. Petrucci |author2=William S. Harwood |author3=Jeffry D. Madura |title=General chemistry: principles and modern applications |url=https://books.google.com/books?id=5z0vAQAAIAAJ |access-date=22 August 2011 |year=2007 |publisher=Pearson/Prentice Hall |isbn=978-0-13-149330-8 |page=691}}</ref>
[[Infrared spectroscopy]] has been used to show that, in solution, dissociation is accompanied by formation of the [[ion pair]] {{H3O+}}·F<sup>−</sup>.<ref name="Giguère"/>
<ref name="Spectral Signatures and Molecular O">{{cite journal |author1=Radu Iftimie |author2=Vibin Thomas |author3=Sylvain Plessis |author4=Patrick Marchand |author5=Patrick Ayotte | title = Spectral Signatures and Molecular Origin of Acid Dissociation Intermediates | journal = [[J. Am. Chem. Soc.]] | doi = 10.1021/ja077846o| year = 2008| volume = 130| pages = 5901–7 | pmid = 18386892 | issue = 18|bibcode=2008JAChS.130.5901I }}</ref>
:{{H2O}} + HF {{Eqm}} {{H3O+}}⋅F<sup>−</sup>{{pad|3em}}[[acid ionization constant|p''K''<sub>a</sub>]] = 3.17
This [[ion pair]] has been characterized in the crystalline state at very low temperature.<ref name=mootz>
{{cite journal |last1=Mootz |first1=D. |title=Crystallochemical Correlate to the Anomaly of Hydrofluoric Acid. |journal=Angew. Chem. Int. Ed. Engl. |date=1981 |volume=20 |issue=123 |page=791 |doi=10.1002/anie.198107911}}</ref> 
Further association has been characterized both in solution and in the solid state.{{citation needed|date=January 2021}}
:HF + F<sup>−</sup> {{Eqm}} {{chem|HF|2|-}}{{pad|3em}}log ''K'' = 0.6

It is assumed that polymerization occurs as the concentration increases. This assumption is supported by the isolation of a salt of a tetrameric anion {{chem|H|3|F|4|-}}<ref>{{cite journal |last1=Bunič |first1=Tina |last2=Tramšek |first2=Melita |last3=Goreshnik |first3=Evgeny |last4=Žemva |first4=Boris |title=Barium trihydrogen tetrafluoride of the composition Ba(H<sub>3</sub>F<sub>4</sub>)<sub>2</sub>: The first example of homoleptic HF metal environment |journal=Solid State Sciences |date=2006 |volume=8 |issue=8 |pages=927–931 |doi=10.1016/j.solidstatesciences.2006.02.045 |bibcode=2006SSSci...8..927B }}</ref> and by low-temperature X-ray crystallography.<ref name=mootz/>  The species that are present in concentrated aqueous solutions of hydrogen fluoride have not all been characterized; in addition to {{chem|HF|2|-}} which is known<ref name="Giguère"/> the formation of other polymeric species, {{chem|H|''n''−1|F|''n''|-}}, is highly likely.

The [[Hammett acidity function]], ''H''<sub>0</sub>, for 100% HF was first reported as −10.2,<ref name="HymanKilpatrick1957">{{cite journal | last1=Hyman | first1=Herbert H. | last2=Kilpatrick | first2=Martin | last3=Katz | first3=Joseph J. | title=The Hammett Acidity Function ''H''<sub>0</sub> for Hydrofluoric Acid Solutions | journal=Journal of the American Chemical Society | publisher=American Chemical Society (ACS) | volume=79 | issue=14 | year=1957 | issn=0002-7863 | doi=10.1021/ja01571a016 | pages=3668–3671| bibcode=1957JAChS..79.3668H }}</ref> while later compilations show −11, comparable to values near −12 for pure [[sulfuric acid]].<ref name="Jolly">{{cite book | last=Jolly | first=William L. | title=Modern Inorganic Chemistry | publisher=McGraw-Hill | year=1984 | isbn=0-07-032768-8 | oclc=22861992 | page=203 }}</ref><ref name="Cotton 109">{{cite book | last1=Cotton | first1=F. A. |author-link = F. Albert Cotton | last2=Wilkinson | first2=G. | title=Advanced Inorganic Chemistry | edition=5th | publisher=Wiley | location=New York | year=1988 | isbn=0-471-84997-9 | oclc=16580057 | page=109 }}</ref>

==Acidity==
Unlike other [[Hydrogen halide|hydrohalic acids]], such as [[hydrochloric acid]], hydrogen fluoride is only a [[weak acid]] in dilute aqueous solution.<ref>{{cite book |last1=Wiberg |first1=Egon |last2=Wiberg |first2=Nils |last3=Holleman |first3=Arnold Frederick |title=Inorganic Chemistry |date=2001 |publisher=Academic Press |isbn=978-0-12-352651-9 |location=San Diego |page=425}}</ref> 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, {{chem|H|2|O}}, and {{chem|F|-}} anions to form clusters.<ref>{{cite web |last=Clark |first=Jim |title=The acidity of the hydrogen halides |url=http://www.chemguide.co.uk/inorganic/group7/acidityhx.html |access-date=4 September 2011 |year=2002}}</ref> At high concentrations, HF molecules undergo [[homoassociation]] to form polyatomic ions (such as [[bifluoride]], {{chem|HF|2|-}}) and [[proton]]s, thus greatly increasing the acidity.<ref name="H+">{{cite book |author1=Chambers, C. |author2=Holliday, A. K. |title=Modern inorganic chemistry (An intermediate text) |year=1975 |publisher=The Butterworth Group |pages=328–329 |url=http://files.rushim.ru/books/neorganika/Chambers.pdf |url-status=dead |archive-url=https://web.archive.org/web/20130323002902/http://files.rushim.ru/books/neorganika/Chambers.pdf |archive-date=2013-03-23 }}</ref> This leads to [[protonation]] of very strong acids like hydrochloric, sulfuric, or nitric acids when using concentrated hydrofluoric acid solutions.<ref>{{cite book |author=Hannan, Henry J. |title=Course in chemistry for IIT-JEE 2011 |year=2010 |publisher=Tata McGraw Hill Education Private Limited |isbn=9780070703360 |pages=15–22 |url=https://books.google.com/books?id=wNYMUeCcaSEC&pg=SA15-PA22}}</ref> Although hydrofluoric acid is regarded as a weak acid, it is very corrosive, even attacking glass when hydrated.<ref name="H+"/>

Dilute solutions are weakly acidic with an [[acid ionization constant]] {{math|size=100%|''K''<sub>a</sub> {{=}} {{val|6.6e-4}}}} (or {{math|size=100%|1=p''K''<sub>a</sub>&nbsp;=&nbsp;3.18}}),<ref name="Pearson/Prentice Hall"/> in contrast to corresponding solutions of the other hydrogen halides, which are strong acids ({{math|size=100%|p''K''<sub>a</sub> < 0}}). However concentrated solutions of hydrogen fluoride are much more strongly acidic than implied by this value, as shown by measurements of the [[Hammett acidity function]] ''H''<sub>0</sub>(or "effective pH"). During self ionization of 100% liquid HF the ''H''<sub>0</sub> was first measured as −10.2<ref name="HymanKilpatrick1957" /> and later compiled as −11, comparable to values near −12 for [[sulfuric acid]].<ref name="Jolly" /><ref name="Cotton 109"/>

In thermodynamic terms, HF solutions are highly [[ideal solution|non-ideal]], with the [[activity (chemistry)|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.<ref>C.&nbsp;E. Housecroft and A.&nbsp;G. Sharpe "Inorganic Chemistry" (Pearson Prentice Hall, 2nd ed. 2005), p. 170.</ref> [[Paul Giguère]] and Sylvia Turrell<ref name="Giguère">{{cite journal
| doi = 10.1021/ja00537a008
| title = The nature of hydrofluoric acid. A spectroscopic study of the proton-transfer complex H<sub>3</sub>O<sup>+</sup>...F<sup>−</sup>
| year = 1980
| author = Giguère, Paul A.  | author2=Turrell, Sylvia
| journal = [[J. Am. Chem. Soc.]]
| volume = 102
| pages = 5473
| issue = 17| author-link = Paul-Antoine Giguère
}}</ref><ref name="Spectral Signatures and Molecular O"/> have shown by [[infrared spectroscopy]] that the predominant solute species in dilute solution is the hydrogen-bonded [[ion pair]] {{H3O+}}·F<sup>−</sup>.<ref name="Cotton 104">{{Harvp|Cotton|Wilkinson|1988|p=104}}</ref>

:{{H2O}} + HF {{Eqm}} {{H3O+}}⋅F<sup>−</sup>

With increasing concentration of HF the concentration of the [[bifluoride|hydrogen difluoride ion]] also increases.<ref name="Giguère"/> The reaction 
:3 HF {{eqm}} {{chem|HF|2|-}} + H<sub>2</sub>F<sup>+</sup>
is an example of  [[homoassociation|homoconjugation]].

==Health and safety==
[[File:61569264 jamesheilman-224x2991.jpg|thumb|upright=1.4|A hydrofluoric acid burn of the hand]]
{{Main|Hydrofluoric acid burn}}
{{More medical citations needed|date=November 2019}}

In addition to being a highly [[corrosive]] liquid, hydrofluoric acid is also a powerful contact [[poison]].  Since it can penetrate tissue, poisoning can occur readily through exposure of skin or eyes, [[inhalation]], or [[ingestion]]. Symptoms of exposure to hydrofluoric acid may not be immediately evident, and this can provide false reassurance to victims, causing them to delay medical treatment.<ref name="pmid11505130">{{cite journal |vauthors=Yamashita M, Yamashita M, Suzuki M, Hirai H, Kajigaya H|title=Ionophoretic delivery of calcium for experimental hydrofluoric acid burns |journal=Crit. Care Med. |volume=29 |issue=8 |pages=1575–8 |year=2001 |pmid=11505130|doi=10.1097/00003246-200108000-00013|s2cid=45595073 }}</ref> Despite its irritating vapor, HF may reach dangerous levels without an obvious odor.<ref name=":0" /> It interferes with nerve function, meaning that burns may not initially be painful. Accidental exposures can go unnoticed, delaying treatment and increasing the extent and seriousness of the injury.<ref name="pmid11505130" /> Symptoms of HF exposure include irritation of the eyes, skin, nose, and throat, eye and skin burns, [[rhinitis]], [[bronchitis]], [[pulmonary edema]] (fluid buildup in the lungs), and bone damage<ref>{{Cite web|title =  NIOSH Pocket Guide to Chemical Hazards – Hydrogen fluoride|url = https://www.cdc.gov/niosh/npg/npgd0334.html|website = CDC |access-date = 2015-11-28}}</ref> due to HF strongly interacting with calcium in bones.<ref>{{cite web |title=Hydrofluoric Acid Fact Sheet |url=https://essr.umd.edu/sites/default/files/2021-10/HydrofluoricAcidFactSheet.pdf |website=Department of Environmental Safety, Sustainability & Risk |publisher=University of Maryland |access-date=11 March 2024 |quote=the HF molecule can cause deep tissue damage, including destruction of the bone. ... when fluoride ions bind to calcium and magnesium}}</ref> In a concentrated form, HF can cause severe tissue destruction through lesions and mucous membrane damage, but dilute HF is still dangerous because of its high lipid affinity, leading to cellular death of nerves, blood vessels, tendons, bones, and other tissues.<ref>{{cite journal |vauthors=Bajraktarova-Valjakova E, Korunoska-Stevkovska V, Georgieva S, Ivanovski K, Bajraktarova-Misevska C, Mijoska A, Grozdanov A|title=Hydrofluoric Acid: Burns and Systemic Toxicity, Protective Measures, Immediate and Hospital Medical Treatment |journal=Open Access Macedonian Journal of Medical Sciences |volume=6 | issue=11| pages=2257–69 | year=2018 |pmid=30559898|doi=10.3889/oamjms.2018.429|doi-broken-date=1 July 2025 |pmc=6290397 }}</ref> Cases of intestinal damage from contact with hydrofluoric acid requiring surgical intervention are rare but known in medical literature.<ref>{{cite journal |last1=Cappell |first1=M. S. |last2=Simon |first2=T. |title=Fulminant acute colitis following a self-administered hydrofluoric acid enema |journal=The American Journal of Gastroenterology |date=January 1993 |volume=88 |issue=1 |pages=122–126 |pmid=8420252 |issn=0002-9270}}</ref>

Hydrofluoric acid burns should be rapidly treated with [[calcium gluconate#Hydrofluoric acid burns|calcium gluconate]] gel.{{fact|date=November 2025}}

==In popular culture==
*In the episodes "[[Cat's in the Bag...]]" and "[[Box Cutter (Breaking Bad)|Box Cutter]]" of the crime drama television series ''[[Breaking Bad]]'', [[Walter White (Breaking Bad)|Walter White]] and [[Jesse Pinkman]] use hydrofluoric acid to chemically degrade  the corpses of gangsters.<ref>{{cite web |url=https://www.bbc.com/news/magazine-23710654 |title=How much of the science in Breaking Bad is real? |date=16 August 2013 |website=BBC News |archive-url=https://web.archive.org/web/20230817013316/https://www.bbc.com/news/magazine-23710654 |archive-date=17 August 2023}}</ref><ref>{{cite web |url=https://edu.rsc.org/analysis/breaking-bad-ii-acid-bath-disposal-of-bodies/3007374.article |title=Breaking Bad II – acid bath disposal of bodies |last=Hare |first=Jonathan |date=1 May 2011 |website=education in chemistry |publisher=Royal Society of Chemistry |archive-url=https://web.archive.org/web/20230610112007/https://edu.rsc.org/analysis/breaking-bad-ii-acid-bath-disposal-of-bodies/3007374.article |archive-date=10 June 2023}}</ref>
* In "Of Past Regret and Future Fear", a [[List of ER episodes|Season 4]] episode of the television medical drama ''[[ER (TV Series)|ER]]'', a patient is brought in with severe chemical burns from hydrofluoric acid that prove to be fatal.
* Hydrofluoric acid is mentioned in the novel ''Virtual Destruction'' by [[Kevin J. Anderson]] and [[Doug Beason]]. A character named Jose Aragon comes in contact with the acid on his hand, prompting him to seek treatment at a medical center.

==See also==
* [[Vapour phase decomposition]]
* [[2019 Philadelphia Energy Solutions refinery explosion]]

==References==
{{reflist}}

==External links==
* [https://web.archive.org/web/20171205143606/https://www.cdc.gov/niosh/ipcsneng/neng0283.html International Chemical Safety Card 0283]
* [https://www.cdc.gov/niosh/npg/npgd0334.html NIOSH Pocket Guide to Chemical Hazards]
* {{PubChemLink|14917}} (HF)
* {{PubChemLink|144681}} (5HF)
* {{PubChemLink|141165}} (6HF)
* {{PubChemLink|144682}} (7HF)
* [http://content.nejm.org/cgi/reprint/356/6/e5.pdf "Hydrofluoric Acid Burn"], ''[[The New England Journal of Medicine]]''—Acid burn case study

{{Authority control}}

[[Category:Fluorides]]
[[Category:Nonmetal halides]]
[[Category:Mineral acids]]
[[Category:Acid catalysts]]