Hydrofluoric acid

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⁻ pK_a = 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, H
_n−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 K_a = 6.6×10⁻⁴ (or pK_a = 3.18), in contrast to corresponding solutions of the other hydrogen halides, which are strong acids (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. 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

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. 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 due to HF strongly interacting with calcium in bones. 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. Cases of intestinal damage from contact with hydrofluoric acid requiring surgical intervention are rare but known in medical literature.

Hydrofluoric acid burns can be treated with calcium gluconate gel. Hexafluorine solution has been claimed to mitigate the adverse effects of burns. However, further studies have found the efficacy of the solution in reducing the longer term effects of electrolyte imbalance is almost identical to that of water rinsing.