STRUCTURE AND BONDING 53force which explains the abnormally high boiling points of hydrogen
fluoride, water and ammonia. The hydrogen bonding in hydrogen
fluoride is so strong that salts of a hypothetical acid H 2 F 2 can be
isolated, for example, KHF 2 with the structure K + [F • • • H • • • F] ~.
Again, ice is known to have a structure similar to that of diamond
with four bonds tetrahedrally arranged. Two hydrogen bonds bind
the lone pairs of electrons on a given oxygen atom to the positively
charged hydrogen atoms of two adjacent water molecules, these
hydrogen bonds being slightly longer than the hydrogen-oxygen
covalent bonds of the water molecule (Figure 2.6).
(a)Figure 2.6. The tetrahedral structures of ice: (a), (b) are planes through sheets of selected
oxygen nuclei (open circles)', hydrogen nuclei (shown in the insert as solid circles) are
not shown in the main drawing. The insert illustrates the overlap of oxygen line pairs
and the hydrogen nuclei, thus forming the hydrogen bonds (dotted lines)
4
The whole structure is rigid but open, giving ice a low density. The
structure of liquid water is similar but less rigid; this explains the
fact that water has a high melting point and dielectric constant (per-
mittivity). Hydrogen bonding has been suggested as one reason why
both H + and OH ~ ions have very high ionic mobilities.
Hydrogen bonding is found between most compounds containing
hydrogen attached to nitrogen, oxygen or fluorine; it explains why,
for example, ethanol C 2 H 5 OH, (C 2 H 6 O) has a boiling point
of 351 K whilst the isomeric dimethyl ether CH 3 —O—CH 3 boils
at 249.4 K, and why some carboxylic acids associate into dimers,
for example ethanoic acid in benzene dimerises to form