negative 'poles' are attracted toward the Na + cation,
and around each Cl - anion is a similar layer with
postive poles aligned towards the negative ion. This
phenomenon of hydration (Box 4.1) is an ion-dipole
interaction . The water molecule possesses no net charge,
but is attracted towards an ion because its oppositely
charged (attracted) pole gets closer to the ion than the
similarly charged (repelled) pole. The attractive force on
the molecule is thus stronger than the repulsive force.
hydrogen is associated with a strongly electronegative
element like oxygen and nitrogen. Hydrogen bonding
is largely responsible for the cohesion between layers in
minerals like kaolinite, for example. It plays a dominant
role in most biological processes, determining for
instance the regular arrangement of polypeptide chains
in protein molecules (Chapter 9), and cross-linking of
the double helix in DNA.
Induced-dipole (van der Waals) interactions
Dipole-dipole interactions: hydrogen bonding
The electrostatic field of a dipole can induce a polariz-
ation in an atom that is not itself intrinsically polarized.
The induced dipole is aligned in such a way that it is
attracted to the original dipole. Curiously this attrac-
tion can operate even between two atoms that have no
permanent dipoles. To see how, we have to look a little
closer at the oscillation of the electron standing wave in
the atom (Chapter 5).
We have gained some understanding of the nature
of electron orbitals in atoms and molecules by compar-
ing them with standing waves on a vibrating string.
Denied by the Uncertainty Principle the opportunity to
follow the electron along a precise course in time, we
have considered the orbital as a time-independent
envelope delineating the electron's spatial domain,
analogous to the wave envelope just visible in a vigor-
ously plucked guitar string. Just as the string actually
oscillates back and forth within its envelope, too fast
for us to see, so the electron must also be vibrating
within its own orbital. In support of this view one may
note that wave mechanics associates an angular
momentum with the trapped electron (in orbitals other
than s-orbitals), suggesting that the electron, in its own
obscure way, does really 'travel round the nucleus'.
The consequence of this oscillation is that if we were
able to freeze an atom or molecule at some instant, we
would find that electron density in each orbital was
not uniformly and symmetrically distributed, but
momentarily concentrated in one part or another, mak-
ing the atom or molecule for that instant a dipole. At
each instant the dipole generates an electric field that
can polarize neighbouring atoms or molecules in con-
cert with its own oscillation, and attract them closer.
Although the electric fields of these synchronous
dipoles average out to zero over a period of time, the
intermolecular attraction they generate does not.
The unique physical and chemical properties of water
(Box 4.1) suggest that some kind of attraction exists
between the molecules in water and ice. This attrac-
tion, known as the hydrogen bond , arises from the elec-
trostatic interaction between one water molecule and
its neighbours: a dipole-dipole interaction . This is seen
most clearly in the regular structure of ice (Figure 7.9c),
in which each hydrogen atom in the molecule is
attracted electrostatically toward the oxygen atom in a
neighbouring molecule. The attraction is maximized if
the geometry allows the hydrogen to associate with
one of the two lone pairs (Figure 7.5c) on the oxygen
atom. Owing to the nearly tetrahedral disposition of
bonding orbitals and lone pairs in the oxygen atom,
this leads not to a close-packed structure but to a three-
dimensional network of molecules analogous to the
structure of diamond (Figure 7.9c, cf. Figure 7.5d).
This ordered structure breaks down during melting,
but about half of the hydrogen bonding is preserved in
liquid water, giving rise to its high viscosity, surface
tension and boiling point (Box 7.7) compared with
other common liquids. Because the molecules are able
to pack closer together in the liquid state than in the
ordered crystalline framework of ice, water exhibits a
slight increase in density on melting, and a correspond-
ing increase in volume on freezing (the reason why ice
floats), another unique property which we take for
granted (Box 2.2).
Hydrogen is the only element capable of 'bridging'
between molecules in this way, a capacity it owes to
the fact that its nucleus is not shielded by any inner
electron shells: it is exposed to the full attraction of
an approaching oxygen lone pair. Hydrogen bonding
has considerable structural significance not only in
water and ice, but in all sorts of compounds in which