Hydrofluoric acid

Last updated
Hydrofluoric acid
Names
IUPAC name
Fluorane [1]
Other names
Fluorhydric acid
Hydronium fluoride
Properties
HF (aq)
AppearanceColorless
Density 1.15 g/mL (for 48% soln.)
Acidity (pKa)3.17 [2]
Hazards [3]
GHS pictograms GHS-pictogram-acid.svg GHS-pictogram-skull.svg
GHS Signal word Danger
H280, H300, H310, H314, H318, H330
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, P410+403, P501
NFPA 704 (fire diamond)
Flammability code 0: Will not burn. E.g. waterHealth code 4: Very short exposure could cause death or major residual injury. E.g. VX gasReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no codeHydrofluoric acid
0
4
0
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

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

Contents

When hydrofluoric acid comes into contact with human skin it causes deep burns.

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 ca. 265 °C.

CaF2 + H2SO4 → 2 HF + CaSO4

The acid is also a by-product of the production of phosphoric acid from apatite/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 plastic containers. [4] [5]

Hydrofluoric acid can be found in nature, having been released in a volcanic eruption.

Properties

In dilute aquous 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 H
3
O+
·F. [11] [12]

H
2
O
+ 2HF H+ + F- + H
3
O+
⋅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. [14]

HF + F- HF2- 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 anion H3F4- [15] and by low-temperature X-ray crystallography. [13] The species that are present in concentrated aqueous solutions of hydrogen fluoride have not been characterized; the formation of polymeric species, Hn-1Fn-, is highly likely.

The Hammett acidity function, H0, for 100% HF is estimated to be between −10.2 and −11. [16] which is comparable to the value −12 for sulfuric acid. [17] [18]

Solutions of hydrofluoric acid attack glass, so they are stored and used in vessels made of teflon. They attack human skin, so must be handled with great care: see #Health and Safety, below.

Acidity

Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in dilute aqueous solution. [19] 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
2
O
, and F
anions to form clusters. [20] At high concentrations, HF molecules undergo homoassociation to form polyatomic ions (such as bifluoride, HF
2
) and protons, thus greatly increasing the acidity. [21] This leads to protonation of very strong acids like hydrochloric, sulfuric, or nitric when using concentrated hydrofluoric acid solutions. [22] Although hydrofluoric acid is regarded as a weak acid, it is very corrosive, even attacking glass when hydrated. [21]

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 Ka = 6.6×104 (or pKa = 3.18), [23] in contrast to corresponding solutions of the other hydrogen halides, which are strong acids (pKa < 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. [24] Paul Giguère and Sylvia Turrell [25] [26] have shown by infrared spectroscopy that the predominant solute species in dilute solution is the hydrogen-bonded ion pair H
3
O+
·F. [27]

H
2
O
+ HF H
3
O+
⋅F

With increasing concentration of HF the concentration of the hydrogen difluoride ion also increases. [25] The reaction

3 HF HF2 + H2F+

is an example of homoconjugation.

Production

Hydrofluoric acid is produced by treatment of the mineral fluorite (CaF2) with concentrated sulfuric acid. When combined at 265 °C, these two substances react to produce hydrogen fluoride and calcium sulfate according to the following chemical equation:

CaF2 + H2SO4 → 2 HF + CaSO4

Although bulk fluorite is a suitable precursor and a major source of world HF production, HF is also produced as a by-product of the production of phosphoric acid, which is derived from the mineral apatite. Apatite sources typically contain a few percent of fluoroapatite, acid digestion of which releases a gaseous stream consisting of sulfur dioxide (from the H2SO4), water, and HF, as well as particulates. After separation from the solids, the gases are treated with sulfuric acid and oleum to afford anhydrous HF. Owing to the corrosive nature of HF, its production is accompanied by the dissolution of silicate minerals, and, in this way, significant amounts of fluorosilicic acid are generated. [4]

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. Because of the ability of hydrofluoric acid to penetrate tissue, poisoning can occur readily through exposure of skin or eyes, or when inhaled or swallowed. 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. [28] Despite having an irritating odor, HF may reach dangerous levels without an obvious odor. [5] HF 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. [28] 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. [29]

In an episode of Breaking Bad titled "Cat's in the Bag...," Walter White uses hydrofluoric acid to dissolve the body of Emilio Koyama. In another episode of Breaking Bad entitled “Box Cutter”, Walter White and Jesse Pinkman use hydrofluoric acid to dissolve the body of Victor.

In the film Saw VI, hydrofluoric acid is used for killing William Easton. In the film Jigsaw Carly is also killed by hydrofluoric acid, injected into her bloodstream.

In an episode of Titans titled "Jason Todd," a young Dick Grayson claims that his parents' murderer used hydrofluoric acid to burn their trapeze ropes.

See also

Related Research Articles

Fluoride is an inorganic, monatomic anion with the chemical formula F
, whose salts are typically white or colorless. Fluoride salts typically have distinctive bitter tastes, and are odorless. Its salts and minerals are important chemical reagents and industrial chemicals, mainly used in the production of hydrogen fluoride for fluorocarbons. Fluoride is classified as a weak base since it only partially associates in solution, but concentrated fluoride is corrosive and can attack the skin.

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

According to the classical definition, a superacid is an acid with an acidity greater than that of 100% pure sulfuric acid, 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 its 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 toxic gas forms white fumes in moist air. It is a useful Lewis acid and a versatile building block for other boron compounds.

A mineral acid is an acid derived from one or more inorganic compounds. All mineral acids form hydrogen ions and the conjugate base when dissolved in water.

Potassium fluoride chemical compound

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 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.

Fluorosulfuric acid (IUPAC name: sulfurofluoridic acid) is the inorganic compound with the chemical formula HSO3F. It is one of the strongest acids commercially available. The formula HSO3F emphasizes its relationship to sulfuric acid, H2SO4; HSO3F is a tetrahedral molecule. It is a colourless liquid, although commercial samples are often yellow.

Cobalt(II) fluoride 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 CoF2·2H2 O and tri-hydrate CoF2·3H2O forms of the compound. The hydrate will also decompose with heat.

Silver(II) fluoride chemical compound (AgF₂)

Silver(II) fluoride is a chemical compound with the formula AgF2. It is a rare example of a silver(II) compound. Silver usually exists in its +1 oxidation state. It is used as a fluorinating agent.

An inorganic nonaqueous solvent is a solvent other than water, that is not an organic compound. Common examples are liquid ammonia, liquid sulfur dioxide, sulfuryl chloride and sulfuryl chloride fluoride, phosphoryl chloride, dinitrogen tetroxide, antimony trichloride, bromine pentafluoride, hydrogen fluoride, pure sulfuric acid and other inorganic acids. These solvents are used in chemical research and industry for reactions that cannot occur in aqueous solutions or require a special environment.

Hydrogen fluoride chemical compound

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. 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.

Fluoroantimonic acid chemical compound

Fluoroantimonic acid is an inorganic compound with the chemical formula H
2
FSbF
6
(also written H
2
F[SbF
6
]
, 2HF·SbF5, or simply HF-SbF5). It is an extremely strong acid, easily qualifying as a superacid. The Hammett acidity function, H0, has been measured for different ratios of HF:SbF5. While the H0 of pure HF is −15, addition of just 1 mol % of SbF5 lowers it to around −20. However, further addition of SbF5 results in rapidly diminishing returns, with the H0 reaching −21 at 10 mol %. The use of an extremely weak base as indicator shows that the lowest attainable H0, even with > 50 mol % SbF5, is somewhere between −21 and −23.

Ammonium bifluoride chemical compound

Ammonium hydrogen fluoride is the inorganic compound with the formula NH4HF2 or NH4F·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 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.

Fluoroboric acid chemical compound

Fluoroboric acid or tetrafluoroboric acid (archaically, fluoboric acid) is an inorganic compound with the chemical formula [H+][BF4-], where H+ represents the solvated proton. The solvent can be any suitably Lewis basic entity. For instance, in water, it can be represented by H
3
OBF
4
(oxonium tetrafluoroborate), although more realistically, several water molecules solvate the proton: [H(H2O)n+][BF4-]. The ethyl ether solvate is also commercially available: [H(Et2O)n+][BF4-], where n is most likely 2. Unlike strong acids like H2SO4 or HClO4, the pure unsolvated substance does not exist (see below).

Potassium bifluoride is the inorganic compound with the formula KHF2. This colourless salt consists of the potassium cation and the bifluoride (HF2) anion. The salt is used in etchant for glass. Sodium bifluoride is related and is also of commercial use as an etchant as well as in cleaning products.

Potassium heptafluorotantalate 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.

Thiophosphoryl fluoride 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.

Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.

Sodium bifluoride chemical compound

Sodium bifluoride is the inorganic compound with the formula NaHF2. It is a salt of sodium cation (Na+) and bifluoride anion (HF2). It is a white, water-soluble solid that decomposes upon heating. Sodium bifluoride is non-flammable, hygroscopic, and has a pungent smell. Sodium bifluoride has a number of applications in industry.

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 5 Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, René; Cuer, Jean Pierre (June 15, 2000), Fluorine Compounds, Inorganic, Weinheim, Germany: Wiley-VCH Verlag GmbH & Co. KGaA, doi:10.1002/14356007.a11_307, ISBN   3-527-30673-0
  5. 1 2 3 4 "CDC – The Emergency Response Safety and Health Database: Systemic Agent: HYDROGEN FLUORIDE/ HYDROFLUORIC ACID – NIOSH". www.cdc.gov. 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. 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. Giguère, Paul A.; Turrell, Sylvia (1980). "The nature of hydrofluoric acid. A spectroscopic study of the proton-transfer complex H
    3
    O+
    ...F". J. Am. Chem. Soc. 102 (17): 5473. doi:10.1021/ja00537a008.
  12. 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. Prkić, Ante; Giljanović, Josipa; Bralić, Marija; Boban, Katarina (2012). "Direct Potentiometric Determination of Fluoride Species by Using Ion-Selective Fluoride Electrode" (PDF). Int. J. Electrochem. Sci. 7: 1170–1179.
  15. 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.
  16. 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. American Chemical Society (ACS). 79 (14): 3668–3671. doi:10.1021/ja01571a016. ISSN   0002-7863.
  17. 1 2 Jolly, William L. (1991). Modern Inorganic Chemistry. McGraw-Hill. p. 203. ISBN   0-07-032768-8. OCLC   22861992.
  18. 1 2 Cotton, F. A.; Wilkinson, G. (1988). Advanced Inorganic Chemistry. New York: Wiley. p. 109. ISBN   0-471-84997-9. OCLC   16580057.
  19. Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic Chemistry. San Diego: Academic Press. p. 425. ISBN   978-0-12-352651-9.
  20. Clark, Jim (2002). "The acidity of the hydrogen halides" . Retrieved 4 September 2011.
  21. 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.
  22. Hannan, Henry J. (2010). Course in chemistry for IIT-JEE 2011. Tata McGraw Hill Education Private Limited. pp. 15–22. ISBN   9780070703360.
  23. 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.
  24. C. E. Housecroft and A. G. Sharpe "Inorganic Chemistry" (Pearson Prentice Hall, 2nd ed. 2005), p. 170.
  25. 1 2 Giguère, Paul A.; Turrell, Sylvia (1980). "The nature of hydrofluoric acid. A spectroscopic study of the proton-transfer complex H
    3
    O+
    ...F". J. Am. Chem. Soc. 102 (17): 5473. doi:10.1021/ja00537a008.
  26. 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.
  27. Cotton & Wilkinson (1988) , p. 104
  28. 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.
  29. "CDC – NIOSH Pocket Guide to Chemical Hazards – Hydrogen fluoride". www.cdc.gov. Retrieved 2015-11-28.