Names | |||
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IUPAC name Fluorane [1] | |||
Other names Fluorhydric acid Hydronium fluoride | |||
Identifiers | |||
3D model (JSmol) | |||
ChEBI | |||
ChemSpider | |||
EC Number |
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PubChem CID | |||
RTECS number |
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UNII | |||
CompTox Dashboard (EPA) | |||
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Properties | |||
HF (aq) | |||
Appearance | Colorless liquid | ||
Density | 1.15 g/mL (for 48% soln.) | ||
Acidity (pKa) | 3.17 [2] | ||
Hazards [3] | |||
GHS labelling: | |||
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) | |||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). |
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.
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]
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]
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]
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]
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. [9] It is now mainly produced by treatment of the mineral fluorite, CaF2, with concentrated sulfuric acid at approximately 265 °C.
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]
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]
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 ]
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]
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]
With increasing concentration of HF the concentration of the hydrogen difluoride ion also increases. [11] The reaction
is an example of homoconjugation.
This article needs more reliable medical references for verification or relies too heavily on primary sources .(November 2019) |
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. [24] 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. [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]
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. [27] [28]
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 (hydroxonium in traditional British English) is the common name for 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 (a positive hydrogen ion, H+) to the surrounding water molecules (H2O). In fact, acids must be surrounded by more than a single water molecule in order to ionize, yielding aqueous H+ and conjugate base. Three main structures for the aqueous proton have garnered experimental support: the Eigen cation, which is a tetrahydrate, H3O+(H2O)3, the Zundel cation, which is a symmetric dihydrate, H+(H2O)2, and the Stoyanov cation, an expanded Zundel cation, which is a hexahydrate: H+(H2O)2(H2O)4. Spectroscopic evidence from well-defined IR spectra overwhelmingly supports the Stoyanov cation as the predominant form. For this reason, it has been suggested that wherever possible, the symbol H+(aq) should be used instead of the hydronium ion.
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.
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.
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.
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.
An inorganic nonaqueous solvent is a solvent other than water, that is not an organic compound. These solvents are used in chemical research and industry for reactions that cannot occur in aqueous solutions or require a special environment. Inorganic nonaqueous solvents can be classified into two groups, protic solvents and aprotic solvents. Early studies on inorganic nonaqueous solvents evaluated ammonia, hydrogen fluoride, sulfuric acid, as well as more specialized solvents, hydrazine, and selenium oxychloride.
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 an aqueous solution termed 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, e.g. 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, much higher than other hydrogen halides.
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.
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.
Ammonium bifluoride is the 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.
Zirconium(IV) fluoride describes members of a family inorganic compounds with the formula (ZrF4(H2O)x. All are colorless, diamagnetic solids. Anhydrous Zirconium(IV) fluoride' is a component of ZBLAN fluoride glass.
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 or tetrafluoroboric acid is an inorganic compound with the simplified chemical formula H+[BF4]−. Unlike other strong acids like H2SO4 or HClO4, the pure tetrafluoroboric acid does not exist. The term "fluoroboric acid" refers to a range of chemical compounds, depending on the solvent. The H+ in the simplified formula of fluoroboric acid represents the solvated proton. The solvent can be any suitable Lewis base. For instance, if the solvent is water, fluoroboric acid can be represented by the formula [H3O]+[BF4]−, although more realistically, several water molecules solvate the proton: [H(H2O)n]+[BF4]−. 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 most likely 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.
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.
Acid strength is the tendency of an acid, symbolised by the chemical formula , to dissociate into a proton, , and an anion, . The dissociation of a strong acid in solution is effectively complete, except in its most concentrated solutions.
Lutetium(III) fluoride is an inorganic compound with a chemical formula LuF3.
the HF molecule can cause deep tissue damage, including destruction of the bone. ... when fluoride ions bind to calcium and magnesium