Hydrogen fluoride is a chemical compound with the chemical formula HF. This colorless gas
or liquid is the principal industrial source of fluorine, often in the aqueous form as
hydrofluoric acid, and thus is the precursor to many important compounds including pharmaceuticals
and polymers. HF is widely used in the petrochemical industry and is a component of many superacids.
Hydrogen fluoride boils near room temperature, while other hydrogen halides evaporate at
much lower temperatures. Unlike other hydrogen halides, HF is lighter than air and diffuses
relatively quickly through porous substances. Hydrogen fluoride is a highly dangerous gas,
forming corrosive and penetrating hydrofluoric acid upon contact with tissue. The gas can
also cause blindness by rapid destruction of the corneas.
French chemist Edmond Frémy is credited with discovering anhydrous hydrogen fluoride while
trying to isolate fluorine, although Carl Wilhelm Scheele prepared hydrofluoric acid
in large quantities in 1771, and this acid was known in the glass industry before then.
Structure
Near or above room temperature, HF is a colorless gas. Below its melting point), HF forms orthorhombic
crystals, consisting of zig-zag chains of HF molecules. The HF molecules, with a short
H–F bond of 95 pm, are linked to neighboring molecules by intermolecular H–F distances
of 155 pm. Liquid HF also consists of chains of HF molecules, but the chains are shorter,
consisting on average of only five or six molecules.
Hydrogen bonding HF molecules interact through hydrogen bonds,
thus creating extra clustering associations with other HF molecules. Because of this,
hydrogen fluoride behaves more like water than like other hydrogen halides, such as
HCl. This hydrogen bonding between HF molecules gives rise to high viscosity in the liquid
phase and lower than expected pressure in the gas phase. Hydrogen fluoride does not
boil until 20 °C in contrast to the heavier hydrogen halides which boil between −85 °C
and −35 °C. Hydrogen fluoride is fully miscible with water,
while the other hydrogen halides have large solubility gaps with water. Hydrogen fluoride
and water also form several compounds in the solid state, most notably a 1:1 compound that
does not melt until −40 °C, which is 44 °C above the melting point of pure HF.
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 other factors such as the tendency
of HF, H 2O, 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 when using concentrated hydrofluoric acid solutions. Although hydrofluoric acid
is regarded as a weak acid, it is very corrosive, even attacking glass when hydrated.
The acidity of hydrofluoric acid solutions vary with concentration owing to hydrogen-bond
interactions of the fluoride ion. Dilute solutions are weakly acidic with an acid ionization
constant Ka = 6.6×10−4, in contrast to corresponding solutions of the other hydrogen
halides which are strong acids. Concentrated solutions of hydrogen fluoride are much more
strongly acid than implied by this value, as shown by measurements of the Hammett acidity
function H0. The H0 for 100% HF is estimated to be between −10.2 and −11, comparable
to the value −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. However, Giguère and Turrell have shown by infrared
spectroscopy that the predominant solute species is the hydrogen-bonded ion-pair [H3O+•F−],
which suggests that the ionization can be described as a pair of successive equilibria:
H2O + HF [H3O+•F−] [H3O+•F−] H3O+ + F−
The first equilibrium lies well to the right and the second to the left, meaning that HF
is extensively dissociated, but that the tight ion pairs reduce the thermodynamic activity
coefficient of H3O+, so that the solution is effectively less acidic.
In concentrated solution, the additional HF causes the ion pair to dissociate with formation
of the hydrogen-bonded hydrogen difluoride ion.
[H3O+•F−] + HF H3O+ + HF2− The increase in free H3O+ due to this reaction
accounts for the rapid increase in acidity, while fluoride ions are stabilized by strong
hydrogen bonding to HF to form HF2−. This interaction between the acid and its own conjugate
base is an example of homoassociation. At the limit of 100% liquid HF, there is self-ionization
3 HF H2F+ + HF2− which forms an extremely acidic solution.
The acidity of anhydrous HF can be increased even further by the addition of Lewis acids
such as SbF5, which can reduce H0 to −21. Solvent
Dry hydrogen fluoride readily dissolves low-valent metal fluorides as well as several molecular
fluorides. Many proteins and carbohydrates can be dissolved in dry HF and recovered from
it. In contrast, most non-fluoride inorganic chemicals react with HF rather than dissolving.
Production and uses Hydrogen fluoride is produced as by the action
of sulfuric acid on pure grades of the mineral fluorite and also as a side-product of the
extraction of the fertilizer precursor phosphoric acid from various minerals. See also hydrofluoric
acid. The anhydrous compound hydrogen fluoride is
more commonly used than its aqueous solution, hydrofluoric acid. HF serves as a catalyst
in alkylation processes in oil refineries. A component of high-octane petrol called "alkylate"
is generated in alkylation units that combine C3 and C4 olefins and iso-butane to generate
petrol. HF is a reactive solvent in the electrochemical
fluorination of organic compounds. In this approach, HF is oxidized in the presence of
a hydrocarbon and the fluorine replaces C–H bonds with C–F bonds. Perfluorinated carboxylic
acids and sulfonic acids are produced in this way.
Hydrogen fluoride is an important catalyst used in the majority of the installed linear
alkyl benzene production in the world. The process involves dehydrogenation of n-paraffins
to olefins, and subsequent reaction with benzene using HF as catalyst.
Elemental fluorine, F2, is prepared by electrolysis of a solution of HF and potassium bifluoride.
The potassium bifluoride is needed because anhydrous hydrogen fluoride does not conduct
electricity. Several million kilograms of F2 are produced annually.
Acyl chlorides or acid anhydrides react with hydrogen fluoride to give acyl fluorides.
HF is often used in palynology to remove silicate minerals, for extraction of dinoflagellate
cysts, acritarchs and chitinozoans. 1,1-Difluoroethane is produced by the mercury-catalyzed
addition of hydrogen fluoride to acetylene: HC≡CH + 2 HF → CH3CHF2
The intermediate in this process is vinyl fluoride, the monomeric percursor to polyvinyl
fluoride. Health effects
Upon contact with moisture, including tissue, hydrogen fluoride immediately converts to
hydrofluoric acid, which is highly corrosive and toxic, and requires immediate medical
attention upon exposure. Breathing in hydrogen fluoride at high levels or in combination
with skin contact can cause death from an irregular heartbeat or from fluid buildup
in the lungs. Notes
References
External links "ATSDR – MMG: Hydrogen Fluoride". Retrieved
May 14, 2006 CDC - NIOSH Pocket Guide to Chemical Hazards