LONDON DISPERSION FORCES
Even though atoms or nonpolar molecules have no dipole moment, they experience attractive
forces among themselves. This is because at any particular point in time, the electron density will be
distributed randomly throughout the orbital; that is, the electron density fluctuates with time. This
leads to rapid polarization and counterpolarization of the electron cloud and thus the formation of
short-lived dipoles. These dipoles interact with the electron clouds of neighboring molecules,
inducing the formation of more dipoles. The attractive interactions of these short-lived dipoles are
called dispersion or London forces.
Dispersion forces are generally weaker than other intermolecular forces. They do not extend over
long distances and are therefore most important when molecules are close together. The strength
of these interactions within a given substance depends directly on how easily the electrons in the
molecules can move (i.e., be polarized). Large molecules in which the electrons are far from the
nucleus are relatively easy to polarize and therefore experience greater dispersion forces among
themselves. If it were not for dispersion forces, the noble gases would not liquefy at any
temperature since no other intermolecular forces exist between the noble gas atoms. The low
temperature at which the noble gases liquefy is indicative of the relatively small magnitude of
dispersion forces between the atoms.
HYDROGEN BONDING
Hydrogen bonding is a specific, unusually strong form of dipole-dipole interaction. When hydrogen
is bound to a highly electronegative atom such as fluorine, oxygen, or nitrogen, the hydrogen atom
carries little of the electron density of the covalent bond, most of which is shifted over to the
electronegative atom. This positively charged hydrogen atom interacts with the partial negative
charge located on the electronegative atoms of nearby molecules, causing the two molecules to
experience an attraction for each other. Substances that display hydrogen bonding tend to have
unusually high boiling points compared with compounds of similar molecular formula that do not
participate in hydrogen bonding. The difference derives from the energy required to break the
hydrogen bonds. Hydrogen bonding is particularly important in the behavior of water, alcohols,
amines, and carboxylic acids. In fact, if it were not for the hydrogen bonding ability of water, life as
we know it would not be possible on Earth.