Hydrofluoric acid

Last updated
Hydrofluoric acid
Hydrogen-fluoride-3D-vdW.png
Water-3D-vdW.png
Fluoride ion2.svg
Oxonium-ion-3D-vdW.png
Hydrogen fluoride.JPG
Names
IUPAC name
Fluorane [1]
Other names
Fluorhydric acid
Hydronium fluoride
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
EC Number
  • 231-634-8
PubChem CID
RTECS number
  • MW7875000
UNII
  • InChI=1S/FH/h1H Yes check.svgY
    Key: KRHYYFGTRYWZRS-UHFFFAOYSA-N Yes check.svgY
  • InChI=1/FH/h1H
    Key: KRHYYFGTRYWZRS-UHFFFAOYAC
  • F.O
  • [F-].[OH3+]
Properties
HF (aq)
AppearanceColorless liquid
Density 1.15 g/mL (for 48% soln.)
Acidity (pKa)3.17 [2]
Hazards [3]
GHS labelling:
GHS-pictogram-acid.svg GHS-pictogram-skull.svg
Danger
H280, 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, P410+P403, P501
NFPA 704 (fire diamond)
NFPA 704.svgHealth 4: Very short exposure could cause death or major residual injury. E.g. VX gasFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazard ACID: Acid
4
0
0
ACID
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Yes check.svgY  verify  (what is  Yes check.svgYX mark.svgN ?)

Hydrofluoric acid is a solution of hydrogen fluoride (HF) in water. Solutions of HF are colorless, acidic and highly 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 fluorine-containing compounds; examples include the commonly used pharmaceutical antidepressant medication fluoxetine (Prozac) and the material PTFE (Teflon). Elemental fluorine is produced from it. It is commonly used to etch glass and silicon wafers.

Contents

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 Teflon, fluoropolymers, fluorocarbons, and refrigerants such as freon. Many pharmaceuticals contain fluorine. [4]

Production of inorganic fluorides

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. [4]

Etchant, cleaner

Wet etching tanks Wet etching tanks at LAAS (6 inches) 0468.jpg
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 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. [5]

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. [6] 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. [5] [7] 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. [8]

Oil refining

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.

Production

Hydrofluoric acid was first prepared in 1771, by Carl Wilhelm Scheele. [9] 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. [4]

Because of its high reactivity toward glass, hydrofluoric acid is stored in fluorinated plastic (often PTFE) containers. [4] [5]

Properties

In dilute aqueous solution hydrogen fluoride behaves as a weak acid, [10] Infrared spectroscopy has been used to show that, in solution, dissociation is accompanied by formation of the ion pair H3O+·F−. [11] [12]

H2O + HF ⇌H3O+⋅F−  pKa = 3.17

This ion pair has been characterized in the crystalline state at very low temperature. [13] Further association has been characterized both in solution and in the solid state.[ citation needed ]

HF + F−⇌HF−
2 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 H
3F−
4 [14] and by low-temperature X-ray crystallography. [13] The species that are present in concentrated aqueous solutions of hydrogen fluoride have not all been characterized; in addition to HF−
2 which is known [11] the formation of other polymeric species, H
n−1F−
n, is highly likely.

The Hammett acidity function, H0, for 100% HF was first reported as -10.2, [15] while later compilations show -11, comparable to values near -12 for pure sulfuric acid. [16] [17]

Acidity

Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in dilute aqueous solution. [18] 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. [19] At high concentrations, HF molecules undergo homoassociation to form polyatomic ions (such as bifluoride, HF−
2) and protons, thus greatly increasing the acidity. [20] This leads to protonation of very strong acids like hydrochloric, sulfuric, or nitric acids when using concentrated hydrofluoric acid solutions. [21] Although hydrofluoric acid is regarded as a weak acid, it is very corrosive, even attacking glass when hydrated. [20]

Dilute solutions are weakly acidic with an acid ionization constant Ka = 6.6×10−4 (or pKa = 3.18), [10] in contrast to corresponding solutions of the other hydrogen halides, which are strong acids (pKa < 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 [15] and later compiled as −11, comparable to values near −12 for sulfuric acid. [16] [17]

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. [22] Paul Giguère and Sylvia Turrell [11] [12] have shown by infrared spectroscopy that the predominant solute species in dilute solution is the hydrogen-bonded ion pair H3O+·F−. [23]

H2O + HF ⇌H3O+⋅F−

With increasing concentration of HF the concentration of the hydrogen difluoride ion also increases. [11] The reaction

3 HF ⇌HF−
2 + H2F+

is an example of homoconjugation.

Health and safety

A hydrofluoric acid burn of the hand 61569264 jamesheilman-224x2991.jpg
A hydrofluoric acid burn of the hand

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. [24] Despite its irritating vapor, HF may reach dangerous levels without an obvious odor. [5] 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. [24] 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 [25] due to HF strongly interacting with calcium in bones. [26] 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. [27]

Hydrofluoric burns are treated with a calcium gluconate gel.

See also

Related Research Articles

<span class="mw-page-title-main">Acid</span> Chemical compound giving a proton or accepting an electron pair

An acid is a molecule or ion capable of either donating a proton (i.e. hydrogen ion, H+), known as a Brønsted–Lowry acid, or forming a covalent bond with an electron pair, known as a Lewis acid.

In chemistry, hydronium is the cation [H3O]+, also written as H3O+, the type of oxonium ion produced by protonation of water. It is often viewed as the positive ion present when an Arrhenius acid is dissolved in water, as Arrhenius acid molecules in solution give up a proton to the surrounding water molecules. In fact, acids must be surrounded by more than a single water molecule in order to ionize, yielding aqueous H+ and conjugate base.

Calcium fluoride is the inorganic compound of the elements calcium and fluorine with the formula CaF2. It is a white solid that is practically insoluble in water. It occurs as the mineral fluorite (also called fluorspar), which is often deeply coloured owing to impurities.

In chemistry, a superacid (according to the original definition) is an acid with an acidity greater than that of 100% pure sulfuric acid (H2SO4), which has a Hammett acidity function (H0) of −12. According to the modern definition, a superacid is a medium in which the chemical potential of the proton is higher than in pure sulfuric acid. Commercially available superacids include trifluoromethanesulfonic acid (CF3SO3H), also known as triflic acid, and fluorosulfuric acid (HSO3F), both of which are about a thousand times stronger (i.e. have more negative H0 values) than sulfuric acid. Most strong superacids are prepared by the combination of a strong Lewis acid and a strong Brønsted acid. A strong superacid of this kind is fluoroantimonic acid. Another group of superacids, the carborane acid group, contains some of the strongest known acids. Finally, when treated with anhydrous acid, zeolites (microporous aluminosilicate minerals) will contain superacidic sites within their pores. These materials are used on massive scale by the petrochemical industry in the upgrading of hydrocarbons to make fuels.

Boron trifluoride is the inorganic compound with the formula BF3. This pungent, colourless, and toxic gas forms white fumes in moist air. It is a useful Lewis acid and a versatile building block for other boron compounds.

<span class="mw-page-title-main">Hydrogen halide</span> Chemical compound consisting of hydrogen bonded to a halogen element

In chemistry, hydrogen halides are diatomic, inorganic compounds that function as Arrhenius acids. The formula is HX where X is one of the halogens: fluorine, chlorine, bromine, iodine, astatine, or tennessine. All known hydrogen halides are gases at standard temperature and pressure.

<span class="mw-page-title-main">Potassium fluoride</span> Ionic compound (KF)

Potassium fluoride is the chemical compound with the formula KF. After hydrogen fluoride, KF is the primary source of the fluoride ion for applications in manufacturing and in chemistry. It is an alkali halide salt and occurs naturally as the rare mineral carobbiite. Solutions of KF will etch glass due to the formation of soluble fluorosilicates, although HF is more effective.

<span class="mw-page-title-main">Fluorosulfuric acid</span> Chemical compound

Fluorosulfuric acid is the inorganic compound with the chemical formula HSO3F. It is one of the strongest acids commercially available. It is a tetrahedral molecule and is closely related to sulfuric acid, H2SO4, substituting a fluorine atom for one of the hydroxyl groups. It is a colourless liquid, although commercial samples are often yellow.

<span class="mw-page-title-main">Cobalt(II) fluoride</span> Chemical compound

Cobalt(II) fluoride is a chemical compound with the formula (CoF2). It is a pink crystalline solid compound which is antiferromagnetic at low temperatures (TN=37.7 K) The formula is given for both the red tetragonal crystal, (CoF2), and the tetrahydrate red orthogonal crystal, (CoF2·4H2O). CoF2 is used in oxygen-sensitive fields, namely metal production. In low concentrations, it has public health uses. CoF2 is sparingly soluble in water. The compound can be dissolved in warm mineral acid, and will decompose in boiling water. Yet the hydrate is water-soluble, especially the di-hydrate CoF2·2H2O and tri-hydrate CoF2·3H2O forms of the compound. The hydrate will also decompose with heat.

<span class="mw-page-title-main">Hydrogen fluoride</span> Chemical compound

Hydrogen fluoride (fluorane) is an inorganic compound with chemical formula HF. 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 polymers such as polytetrafluoroethylene (PTFE). HF is also widely used in the petrochemical industry as a component of superacids. Due to strong and extensive hydrogen bonding, it boils at near room temperature, which is much higher of a temperature than other hydrogen halides.

<span class="mw-page-title-main">Fluoroantimonic acid</span> Chemical compound

Fluoroantimonic acid is a mixture of hydrogen fluoride and antimony penta­fluoride, containing various cations and anions. This mixture is a superacid that, in terms of corrosiveness, is trillions of times stronger than pure sulfuric acid when measured by its Hammett acidity function. It even protonates some hydro­carbons to afford pentacoordinate carbo­cations. Like its precursor hydrogen fluoride, it attacks glass, but can be stored in containers lined with PTFE (Teflon) or PFA.

<span class="mw-page-title-main">Hexafluorosilicic acid</span> Octahedric silicon compound

Hexafluorosilicic acid is an inorganic compound with the chemical formula H
2SiF
6. Aqueous solutions of hexafluorosilicic acid consist of salts of the cation and hexafluorosilicate anion. These salts and their aqueous solutions are colorless.

<span class="mw-page-title-main">Ammonium bifluoride</span> Chemical compound

Ammonium bifluoride is an inorganic compound with the formula [NH4][HF2] or [NH4]F·HF. It is produced from ammonia and hydrogen fluoride. This colourless salt is a glass-etchant and an intermediate in a once-contemplated route to hydrofluoric acid.

The bifluoride ion is an inorganic anion with the chemical formula [HF2]−. The anion is colorless. Salts of bifluoride are commonly encountered in the reactions of fluoride salts with hydrofluoric acid. The commercial production of fluorine involves electrolysis of bifluoride salts.

The Hammett acidity function (H0) is a measure of acidity that is used for very concentrated solutions of strong acids, including superacids. It was proposed by the physical organic chemist Louis Plack Hammett and is the best-known acidity function used to extend the measure of Brønsted–Lowry acidity beyond the dilute aqueous solutions for which the pH scale is useful.

<span class="mw-page-title-main">Fluoroboric acid</span> Chemical compound

Fluoroboric acid or tetrafluoroboric acid is an inorganic compound with the simplified chemical formula H+[BF4]−. Solvent-free tetrafluoroboric acid has not been reported. The term "fluoroboric acid" usually refers to a range of compounds including hydronium tetrafluoroborate, which are available as solutions. The ethyl ether solvate is also commercially available, where the fluoroboric acid can be represented by the formula [H( 2O)n]+[BF4]−, where n is 2.

An acidity function is a measure of the acidity of a medium or solvent system, usually expressed in terms of its ability to donate protons to a solute. The pH scale is by far the most commonly used acidity function, and is ideal for dilute aqueous solutions. Other acidity functions have been proposed for different environments, most notably the Hammett acidity function, H0, for superacid media and its modified version H− for superbasic media. The term acidity function is also used for measurements made on basic systems, and the term basicity function is uncommon.

<span class="mw-page-title-main">Potassium heptafluorotantalate</span> Chemical compound

Potassium heptafluorotantalate is an inorganic compound with the formula K2[TaF7]. It is the potassium salt of the heptafluorotantalate anion [TaF7]2−. This white, water-soluble solid is an intermediate in the purification of tantalum from its ores and is the precursor to the metal.

<span class="mw-page-title-main">Thiophosphoryl fluoride</span> Chemical compound

Thiophosphoryl fluoride is an inorganic molecular gas with formula PSF3 containing phosphorus, sulfur and fluorine. It spontaneously ignites in air and burns with a cool flame. The discoverers were able to have flames around their hands without discomfort, and called it "probably one of the coldest flames known". The gas was discovered in 1888.

<span class="mw-page-title-main">Lutetium(III) fluoride</span> Chemical compound

Lutetium(III) fluoride is an inorganic compound with a chemical formula LuF3.

References

  1. ↑ Favre, Henri A.; Powell, Warren H., eds. (2014). Nomenclature of Organic Chemistry: IUPAC Recommendations and Preferred Names 2013. Cambridge: The Royal Society of Chemistry. p. 131. ISBN   9781849733069.
  2. ↑ Harris, Daniel C. (2010). Quantitative Chemical Analysis (8th international ed.). New York: W. H. Freeman. pp. AP14. ISBN   978-1429263092.
  3. ↑ "Hydrofluoric Acid". PubChem. National Institute of Health. Retrieved October 12, 2017.
  4. 1 2 3 4 Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, René; Cuer, Jean Pierre (2000). "Fluorine Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry . Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_307. ISBN   3527306730.
  5. 1 2 3 4 "Hydrogen Fluoride/Hydrofluoric Acid: Systemic Agent". Emergency Response Safety and Health Database. NIOSH - CDC. May 12, 2011. Archived from the original on Dec 7, 2015. Retrieved 2015-12-04.
  6. ↑ Craig, Robert (2006). Craig's restorative dental materials. St. Louis, Mo: Mosby Elsevier. ISBN   0-323-03606-6. OCLC   68207297.
  7. ↑ Strachan, John (January 1999). "A deadly rinse: The dangers of hydrofluoric acid". Professional Carwashing & Detailing. 23 (1). Archived from the original on April 25, 2012.
  8. ↑ Edwards, D. (1982). "Fragmentary non-vascular plant microfossils from the late Silurian of Wales". Botanical Journal of the Linnean Society. 84 (3): 223–256. doi:10.1111/j.1095-8339.1982.tb00536.x.
  9. ↑ Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. p. 921. ISBN   978-0-08-022057-4.
  10. 1 2 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.
  11. 1 2 3 4 Giguère, Paul A.; Turrell, Sylvia (1980). "The nature of hydrofluoric acid. A spectroscopic study of the proton-transfer complex H3O+...F−". J. Am. Chem. Soc. 102 (17): 5473. doi:10.1021/ja00537a008.
  12. 1 2 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.
  13. 1 2 Mootz, D. (1981). "Crystallochemical Correlate to the Anomaly of Hydrofluoric Acid". Angew. Chem. Int. Ed. Engl. 20 (123): 791. doi:10.1002/anie.198107911.
  14. ↑ Bunič, Tina; Tramšek, Melita; Goreshnik, Evgeny; Žemva, Boris (2006). "Barium trihydrogen tetrafluoride of the composition Ba(H3F4)2: The first example of homoleptic HF metal environment". Solid State Sciences. 8 (8): 927–931. Bibcode:2006SSSci...8..927B. doi:10.1016/j.solidstatesciences.2006.02.045.
  15. 1 2 Hyman, Herbert H.; Kilpatrick, Martin; Katz, Joseph J. (1957). "The Hammett Acidity Function H0 for Hydrofluoric Acid Solutions". Journal of the American Chemical Society. 79 (14). American Chemical Society (ACS): 3668–3671. doi:10.1021/ja01571a016. ISSN   0002-7863.
  16. 1 2 Jolly, William L. (1984). Modern Inorganic Chemistry. McGraw-Hill. p. 203. ISBN   0-07-032768-8. OCLC   22861992.
  17. 1 2 Cotton, F. A.; Wilkinson, G. (1988). Advanced Inorganic Chemistry (5th ed.). New York: Wiley. p. 109. ISBN   0-471-84997-9. OCLC   16580057.
  18. ↑ Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic Chemistry. San Diego: Academic Press. p. 425. ISBN   978-0-12-352651-9.
  19. ↑ Clark, Jim (2002). "The acidity of the hydrogen halides" . Retrieved 4 September 2011.
  20. 1 2 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.
  21. ↑ Hannan, Henry J. (2010). Course in chemistry for IIT-JEE 2011. Tata McGraw Hill Education Private Limited. pp. 15–22. ISBN   9780070703360.
  22. ↑ C. E. Housecroft and A. G. Sharpe "Inorganic Chemistry" (Pearson Prentice Hall, 2nd ed. 2005), p. 170.
  23. ↑ Cotton & Wilkinson (1988) , p. 104
  24. 1 2 Yamashita M, Yamashita M, Suzuki M, Hirai H, Kajigaya H (2001). "Ionophoretic delivery of calcium for experimental hydrofluoric acid burns". Crit. Care Med. 29 (8): 1575–8. doi:10.1097/00003246-200108000-00013. PMID   11505130. S2CID   45595073.
  25. ↑ "NIOSH Pocket Guide to Chemical Hazards – Hydrogen fluoride". CDC. Retrieved 2015-11-28.
  26. ↑ "Hydrofluoric Acid Fact Sheet" (PDF). Department of Environmental Safety, Sustainability & Risk. University of Maryland. Retrieved 11 March 2024. the HF molecule can cause deep tissue damage, including destruction of the bone. ... when fluoride ions bind to calcium and magnesium
  27. ↑ Bajraktarova-Valjakova E, Korunoska-Stevkovska V, Georgieva S, Ivanovski K, Bajraktarova-Misevska C, Mijoska A, Grozdanov A (2018). "Hydrofluoric Acid: Burns and Systemic Toxicity, Protective Measures, Immediate and Hospital Medical Treatment". Open Access Macedonian Journal of Medical Sciences. 6 (11): 2257–69. doi:10.3889/oamjms.2018.429 (inactive 2024-10-11). PMC   6290397 . PMID   30559898.{{cite journal}}: CS1 maint: DOI inactive as of October 2024 (link)
  28. ↑ "How much of the science in Breaking Bad is real?". BBC News. 16 August 2013. Archived from the original on 17 August 2023.
  29. ↑ Hare, Jonathan (1 May 2011). "Breaking Bad II – acid bath disposal of bodies". education in chemistry. Royal Society of Chemistry. Archived from the original on 10 June 2023.
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