This change in properties cannot be simply accounted for in terms
of bond energies; the mean X—H bond energy increases from
nitrogen to fluorine, and hydrogen fluoride has a large bonddissociation
energy (566kJmol~1). But we note that in the CH4
molecule there are no lone pairs of electrons—all four valency
electrons are involved in bonding. In ammonia, there is one lone
pair, which as we have seen can be donated either to a proton
(making ammonia a Lowry-Br0nsted base, NH3 + H + ^NH^)
or to another acceptor molecule (making ammonia a Lewis base,
p. 91). The molecules H2O and HF have two and three lone pairs
respectively; falling-off of base strength implies that the presence of
more than one lone pair reduces the donor power of the molecule.
But, obviously, the appearance of acidic behaviour implies that the
bond X—H is more readily broken heterolytically i.e. to give X~ +
H +. We may ascribe this to polarity of the bond, i.e. by saying that
the pair of electrons in the covalent H—F bond is closer to the
fluorine than to the hydrogen. Unfortunately, there is no very sure
method of ascertaining this bond polarity (the fact that hydrogen
fluoride HF has a dipole moment means that the molecule as a
+ —
whole is polar in, presumably, the sense H—F, but this does not
necessarily tell us about the bond polarity). Another way of describing
this trend towards acidity is to say that the electronegativity
of the element increases from carbon to fluorine. We may simply
note that this trend to acidity is also apparent in other periods, for
example, in Period 3, silane SiH4 is non-acidic and non-basic.