hydrofluoric acid

The principal use of hydrofluoric acid is in organofluorine chemistry. Many organofluorine compounds are prepared using HF as the fluorine source, including Teflon, fluoropolymers, fluorocarbons, and refrigerants such as freon. Many pharmaceuticals contain fluorine.

Most high-volume inorganic fluoride compounds are prepared from hydrofluoric acid. Foremost are Na₃AlF₆, cryolite, and AlF₃, 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.

File:Wet etching tanks at LAAS (6 inches) 0468.jpg

It is used in the semiconductor industry as a major component of Wright etch and buffered oxide etch, which are used to clean silicon wafers. In a similar manner it is also used to 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.

:SiO₂ + 4 HF → SiF₄(g) + 2 H₂O

:SiO₂ + 6 HF → H₂SiF₆ + 2 H₂O

A 5% to 9% hydrofluoric acid gel is also commonly used to etch all ceramic dental restorations to improve bonding.{{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}} For similar reasons, dilute hydrofluoric acid is a component of household rust stain remover, in car washes in "wheel cleaner" compounds, in ceramic and fabric rust inhibitors, and in water spot removers.{{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 }} 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 macerations 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.{{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

}}

In a standard oil refinery process known as alkylation, isobutane is alkylated with low-molecular-weight alkenes (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 carbocations, which alkylate isobutane. The reaction is carried out at mild temperatures (0 and 30 °C) in a two-phase reaction.

Hydrofluoric acid was first prepared in 1771, by Carl Wilhelm Scheele.{{Greenwood&Earnshaw1st|page=921}} It is now mainly produced by treatment of the mineral fluorite, CaF₂, with concentrated sulfuric acid at approximately 265 °C.

:CaF₂ + H₂SO₄ → 2 HF + CaSO₄

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.{{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}}

Because of its high reactivity toward glass, hydrofluoric acid is stored in fluorinated plastic (often PTFE) containers.{{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 }}

In dilute aqueous solution hydrogen fluoride behaves as a weak acid,{{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}}

Infrared spectroscopy has been used to show that, in solution, dissociation is accompanied by formation of the ion pair {{H3O+}}·F⁻.

{{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}}

:{{H2O}} + HF {{Eqm}} {{H3O+}}⋅F⁻{{pad|3em}}pK_a = 3.17

This ion pair has been characterized in the crystalline state at very low temperature.

{{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}}

Further association has been characterized both in solution and in the solid state.{{citation needed|date=January 2021}}

:HF + F⁻ {{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|-}}{{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(H₃F₄)₂: 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 }} and by low-temperature X-ray crystallography. 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 the formation of other polymeric species, {{chem|H|n−1|F|n|-}}, is highly likely.

The Hammett acidity function, H₀, for 100% HF was first reported as −10.2,{{cite journal | last1=Hyman | first1=Herbert H. | last2=Kilpatrick | first2=Martin | last3=Katz | first3=Joseph J. | title=The Hammett Acidity Function H₀ 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}} while later compilations show −11, comparable to values near −12 for pure sulfuric acid.{{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 }}{{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 }}

Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in dilute aqueous solution.{{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}} 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.{{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}} At high concentrations, HF molecules undergo homoassociation to form polyatomic ions (such as bifluoride, {{chem|HF|2|-}}) and protons, thus greatly increasing the acidity.{{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 }} This leads to protonation of very strong acids like hydrochloric, sulfuric, or nitric acids when using concentrated hydrofluoric acid solutions.{{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}} Although hydrofluoric acid is regarded as a weak acid, it is very corrosive, even attacking glass when hydrated.

Dilute solutions are weakly acidic with an acid ionization constant {{math|size=100%|K_a {{=}} {{val|6.6e-4}}}} (or {{math|size=100%|1=pK_a = 3.18}}), in contrast to corresponding solutions of the other hydrogen halides, which are strong acids ({{math|size=100%|pK_a < 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 H₀(or "effective pH"). During self ionization of 100% liquid HF the H₀ was first measured as −10.2 and later compiled as −11, comparable to values near −12 for sulfuric acid.

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.C. E. Housecroft and A. G. Sharpe "Inorganic Chemistry" (Pearson Prentice Hall, 2nd ed. 2005), p. 170. Paul Giguère and Sylvia Turrell{{cite journal

| doi = 10.1021/ja00537a008

| title = The nature of hydrofluoric acid. A spectroscopic study of the proton-transfer complex H₃O⁺...F⁻

| 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

}} have shown by infrared spectroscopy that the predominant solute species in dilute solution is the hydrogen-bonded ion pair {{H3O+}}·F⁻.{{Harvp|Cotton|Wilkinson|1988|p=104}}

:{{H2O}} + HF {{Eqm}} {{H3O+}}⋅F⁻

With increasing concentration of HF the concentration of the hydrogen difluoride ion also increases. The reaction

:3 HF {{eqm}} {{chem|HF|2|-}} + H₂F⁺

is an example of homoconjugation.

File:61569264 jamesheilman-224x2991.jpg

{{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.{{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 }} Despite its irritating vapor, HF may reach dangerous levels without an obvious odor. 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. 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{{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}} due to HF strongly interacting with calcium in bones.{{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}} 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.{{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 November 2024 |pmc=6290397 }}

Hydrofluoric burns are treated with a calcium gluconate gel.

• In the episodes "Cat's in the Bag..." and "Box Cutter" of the crime drama television series Breaking Bad, Walter White and Jesse Pinkman use hydrofluoric acid to chemically disincorporate bodies of gangsters.{{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}}{{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}}
• In Season 4 episode 20 "Of Past Regret and Future Fear" of the television medical drama 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, and he furiously seeks treatment at a medical center.

• Vapour phase decomposition
• 2019 Philadelphia Energy Solutions refinery explosion

{{reflist}}

• [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

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Category:Fluorides

Category:Nonmetal halides

Category:Mineral acids

Category:Acid catalysts