Biomedical Engineering Reference
In-Depth Information
3
Balls on Springs
The theory of intermolecular forces relates to atomic and/or molecular species that are some
distance apart (say, a few bond lengths). We saw in Chapter 2 that progress can be made in
such a theory without normally invoking the concepts of quantum mechanics. If we want
to truly understand why two atoms combine together to give a chemical bond, and how
bonds get broken and reformed in chemical reactions, then we enter the realms of
valence
theory
. Quantum mechanics plays a dominant part in such discussions.
These are simple-minded comments and my arbitrary division of molecular interactions
is subjective. At first sight, the stability of an NaCl ion pair can be explained in terms
of elementary electrostatics, and we can usefully model argon liquid without recourse to
quantummechanics (apart from the London dispersion potential, which is a 'pure' quantum
mechanical effect). AC-C bond in ethane is at first sight a quantummechanical animal, and
we will certainly have to invoke quantummechanical ideas to explain the reaction of ethene
with dichlorine. But there are grey areas that I can bring to your attention by considering the
phenomenon of hydrogen bonding. The hydrogen bond is an attractive interaction between
a proton donor X-H and a proton acceptor Y in the same or a different molecule
X
−
H ...Y
The bond is usually symbolized by three dots as shown above, in order to reconcile the
existence of compounds such as
NH
3
...HCl
with the trivalence of nitrogen, the divalence of oxygen in oxonium salts and other com-
pounds that apparently break the classical valence rules. Hydrogen bonds typically have
strengths of 10-100 kJ mol
−
1
. The lone pairs of oxygen and nitrogen and the partially
charged character of the proton were eventually recognized as the sources of this bond. The
first reference to this 'weak bond' was made by Latimer and Rodebush (1920).
