Page 33 - Materials Chemistry, Second Edition

20 2 Solid-State Chemistry

materials, the melting points vary significantly depending on the nature of interactions
among the polymer subunits. As Chapter 5 will delineate, these interactions also
greatly affect many other physical properties of these materials.
Molecular solids may exhibit either crystalline or amorphous structures, depend-
ing on the complexity of the individual molecules comprising the bulk material.
As with all solids, the more complex the subunits are, the harder it is for them to
organize themselves in a repeatable fashion, resulting in an amorphous structure.
Unlike purely ionic solids, molecular compounds may be soluble in either nonpolar
or polar solvents, as long as the solvent polarity between solute and solvent is
matched (“like dissolves like”).
Both dipole–dipole and London Dispersion forces are subclasses of van der Waal
interactions. When two polar molecules approach one another, a natural attraction
known as dipole–dipole forces is created between oppositely charged ends.
The relative intensity of dipole–dipole forces may be represented by Eq. 2:
2 2
2 m m 1

ð2Þ 1 2 6
kTr
3 4pe o
where m 1 and m 2 are the molecular dipole moments (Debyes); r, the average distance
˚
of separation (A); T, the temperature (K); and k is the Boltzmann constant
1
(1.38065 10 23 JK ).
In addition to the mutual attraction between polar molecules, there may also be an
interaction between the solute molecules and liquid or gaseous solvent. In highly
polar solvents such as water or alcohols, a dense shell of solvent molecules will
surround the polar molecules. Although this solute/solvent interaction assists in the
solubility of the molecules in the solvent, the dipole–dipole interactions between
individual molecules are suppressed.
In contrast to dipole–dipole forces, London Dispersion interactions are much
weaker in nature since they involve nonpolar molecules that do not possess permanent
dipole moments. The only modes for molecular attraction are through polarization of
electrons, which leads to the creation of small dipole–dipole interactions and mutual
attractive forces. Since electron polarization occurs much more readily for electrons
farther from the nucleus, this effect is more pronounced for molecules that are larger
with a greater number of electrons, especially positioned on atoms with a high atomic
number, consisting of more diffuse orbitals. These “induced dipole” forces are
responsible for the liquefaction of gases such as He and Ar at low temperatures and
pressures. The relative strength of London Dispersion forces is described by Eq. 3:
3 I 1 I 2 a 1 a 2
ð3Þ 6
2 I 1 þ I 2 r
where I 1 and I 2 are the ionization potentials of the molecules; and a 1 and a 2 are the
polarizabilities of the molecules.
If both polar and nonpolar molecules are present, a dipole-induced dipole inter-
action may occur. For this situation, the strength of association may be represented
by Eq. 4, which is dependent on both the dipole moment of the polar molecule, and