Box 7.7 Estimating the strength of chemical bonds
the strength of a chemical bond is expressed in terms of
the energy required to break it. this is usually quoted in
the form of the molar enthalpy change for a specific hypo-
thetical dissociation reaction such as:
thermal vibration (which is roughly 2 RT per mole) exceeds
the energy of cohesion holding the molecules together.
elements and compounds of low molecular weight that
are gaseous at room temperature (he, ar, CO 2 , N 2 ) have
weak intermolecular forces, but in a solid or liquid material
such as water the forces of cohesion are quite strong.
this enthalpy change, known as the dissociation energy
of molecular hydrogen, indicates how much energy (in
kJ mol -1 ) is required to split every molecule in one mole of
hydrogen gas into two separate atoms. Some relevant dis-
sociation energies are given in table 7.7.1.
Detailed interpretation of these numbers (for example
the estimation of the energy of an individual C-C bond) is
fraught with pitfalls. We shall simply note the following
Table 7.7.1 energies of various types of chemical
Type of interaction
Energy/kJ · mol -1
(i) Covalent and ionic bond energies fall within a similar
range. energies of single covalent bonds are generally
between 200 and 500 kJ mol -1 . Note the similarity to
the activation energies noted for silicate reactions in
Chapter 3 (Figures 3.6 and 3.8).
(ii) Metals have lower values, those for soft metals being
lower than those for hard metals (Na < pb < Ni).
(iii) the energy of the hydrogen bond is roughly a factor
of ten less than for covalent or ionic bonds. For an
individual hydrogen bond the energy is around
24 kJ mol -1 .
(iv) Van der Waals energies are an order of magnitude
lower again than hydrogen bonds.
Covalent bond CC
466 (x 2)
In the absence of quantitative data, one can make a
crude estimate of the energy of interaction between mol-
ecules in a covalent substance from its boiling point. this
is the temperature at which the energy of the molecules'
Van de Waals
The induced-dipole or van der Waals interaction
(named after Dutch physicist and Nobel laureate
Johannes van der Waals) is responsible for a weak
attraction operating between any pair of atoms or
molecules that are sufficiently close (a few tenths of
a nanometre), although its effect can be detected
only when other inter-atomic forces are absent. Van
der Waals forces are responsible for holding the
sheets together in crystals of graphite, and no doubt
contribute to the cohesion of soft sheet silicates like
pyrophyllite and talc. The low hardness of graphite
(1-2) and the large interlayer distance (Box 7.4) pro-
vide a measure of how feeble the attraction is. The
'bond energy' for the van der Waals interaction is typ-
ically two orders of magnitude less than for ionic and
covalent bonding (Box 7.7). The interaction explains
why all substances, even noble gases, form crystals at
sufficiently low temperatures (neon, for instance, melts