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 stronger solutions are available. In industry, it is less important than its anhydrous derivative hydrogen fluoride but in the laboratory it is a more common reagent used to make some organofluorine compound and other fluorides.
Uses
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.
Etchant, cleaner
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 used to etch glass by treatment with silicon dioxide to form gaseous or water-soluble silicon fluorides. It can 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 commonly used to etch all ceramic dental restorations to improve bonding. 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. 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 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.
Production Hydrofluoric acid was first prepared in 1771, by Carl Wilhelm Scheele. 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. Because of its high reactivity toward glass, hydrofluoric acid is stored in fluorinated plastic (often PTFE) containers.
Properties In dilute aqueous solution, hydrogen fluoride behaves as a weak acid. Infrared spectroscopy has been used to show that, in solution, dissociation is accompanied by formation of the ion pair H3O+·F−.
H2O + HF ⇌ H3O+⋅F− pKa = 3.17 This ion pair has been characterized in the crystalline state at very low temperature. Further association has been characterized both in solution and in the solid state.
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 H3F−4 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 HF−2 which is known the formation of other polymeric species, Hn−1F−n, is highly likely. The Hammett acidity function, H0, for 100% HF was first reported as −10.2, while later compilations show −11, comparable to values near −12 for pure sulfuric acid.
Acidity Unlike other hydrohalic acids, such as hydrochloric acid, hydrogen fluoride is only a weak acid in dilute aqueous solution. 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, H2O, and F− anions to form clusters. At high concentrations, HF molecules undergo homoassociation to form polyatomic ions (such as bifluoride, HF−2) and protons, thus greatly increasing the acidity. This leads to protonation of very strong acids like hydrochloric, sulfuric, or nitric acids when using concentrated hydrofluoric acid solutions. 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 Ka = 6.6×10−4 (or pKa = 3.18), 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 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. Paul Giguère and Sylvia Turrell have shown by infrared spectroscopy that the predominant solute species in dilute solution is the hydrogen-bonded ion pair H3O+·F−.
H2O + HF ⇌ H3O+·F− With increasing concentration of HF the concentration of the hydrogen difluoride ion also increases. The reaction
3 HF ⇌ HF−2 + H2F+ is an example of homoconjugation.
Health and safety
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