Trends in physical and chemical properties of hydrogen halides

Preparation of the hydrogen halides: direct combination of hydrogen with a halogen is a general (though not always practical) route: hydrogen and fluorine react explosively/violently, while hydrogen with bromine or iodine reacts reversibly and does not give a pure product directly. A more general route uses displacement: concentrated sulphuric acid displaces hydrogen chloride and (at higher temperature, via the hydrogen sulphate intermediate) hydrogen fluoride from the corresponding ionic halide salts; however, hydrogen bromide and hydrogen iodide cannot be prepared this way, since concentrated sulphuric acid is a strong enough oxidising agent to oxidise HBr and especially HI as fast as they form. Instead, gaseous hydrogen halides (except hydrogen fluoride) can be obtained by adding water dropwise to a phosphorus trihalide (PX3 + 3H2O -> H3PO3 + 3HX); in practice, hydrogen bromide is made by adding bromine dropwise to a paste of red phosphorus and water, and hydrogen iodide by adding water dropwise to a mixture of red phosphorus and iodine (2P + 3X2 -> 2PX3, then 2PX3 + 3H2O -> 2H3PO3 + 6HX, where X = Br or I); any escaping halogen vapour is removed by passing the gas through a column of moist red phosphorus. Hydrogen halides can also be obtained from covalent hydrides: hydrogen sulphide reduces the halogens to the corresponding hydrogen halide (H2S + X2 -> 2HX + S), and hydrogen chloride is also generated as a useful by-product from reactions between hydrocarbons and chlorine.

General properties and trends: the bond dissociation enthalpy falls steadily from HF (+562 kJ/mol) to HCl (+431), HBr (+366) and HI (+299), and the percentage ionic character of the bond likewise falls sharply from 43% in HF to just 7% in HI, reflecting the decreasing electronegativity difference between hydrogen and the halogen down the group. In line with the decreasing bond dissociation enthalpy, the thermal stability of the hydrogen halides falls from HF to HI: for instance, hydrogen iodide decomposes noticeably at 400 °C, whereas hydrogen fluoride and hydrogen chloride remain stable at that temperature. At room temperature all four are gases, but hydrogen fluoride is exceptionally easy to liquefy (owing to strong hydrogen bonding, absent in the others, which also gives HF an anomalously high melting and boiling point compared with the general downward-then-upward trend expected purely from molecular mass); moist air causes all of them to fume, from droplets of the corresponding hydrohalic acid.

Acidic properties: the hydrogen halides are extremely soluble in water because of extensive ionisation (HX + H2O -> H3O+ + X-, X = F, Cl, Br or I), and their aqueous solutions are known as hydrohalic acids. Hydrochloric, hydrobromic and hydroiodic acids are all almost completely ionised in water and are therefore strong acids, but hydrofluoric acid is comparatively weak — a 0.1 mM solution is only about 10% ionised — though it becomes a markedly stronger acid at high concentration (5 M-15 M) because of a secondary equilibrium in which fluoride ion associates with undissociated HF to form the bifluoride ion (HF + H2O <=> H3O+ + F-, and HF + F- <=> HF2-); removal of free fluoride ion by this second equilibrium pulls the first equilibrium further to the right, increasing the effective hydrogen-ion concentration. Several stable salts of the bifluoride ion are known — NaHF2, KHF2 and NH4HF2 — a type of salt the other hydrogen halides do not form. Hydrohalic acids display the typical properties of acids, forming salts with bases and liberating hydrogen with reactive metals; notably, moist (though not dry) hydrofluoric acid rapidly attacks silica and glass (SiO2 + 4HF -> SiF4 + 2H2O; Na2SiO3 + 6HF -> Na2SiF6 + 3H2O), a property unique among the hydrohalic acids and widely used for glass etching. …