Write an essay on intermolecular bonding. Explain how each type of bond arises
and the evidence for the existence of each. Comment on their strengths in
relation to the types of atoms involved; the covalent bond and relative to each
other. Use the concepts of different types and strengths of intermolecular bonds
to explain the following:

There exists four types of intermolecular bonding, they include ionic, covalent,
Van der waals and hydrogen bonding. In order to describe the existence of such
bonding you must also understand the concepts of polarity, polar and non-polar,
and electronegativity.

Ionic bonds are created by the complete transfer of electrons from one atom to
another. In this process of electron transfer, each atom becomes a ion that is
isoelectronic with the nearest noble gas., the substance is held together by
electrostatic forces between the ions. The tendency for these ions to be formed
by elements is corespondent to the octet rule, when atoms react,, they tend to
do so in such a way that they attain an outer shell containing eight electrons.
The factors that effect the formation of ions are ionization energy, electron
affinity, lattice energy.

Figure 1

The transfer of electrons involved in the formation of (a) sodium chloride and
(b) calcium fluoride. Each atom forms an ion with an outer shell containing
eight electrons.

For many elements, compounds cannot be formed by the production of ions, since
the energy released in the formation of the lattice of ions would be
insufficient to overcome the energy required to form the ions would be
insufficient to overcome the energy required to form the ions in the first place.
In order for the atoms to achieve a noble gas configuration they must use
another method of bonding by the process of electron sharing. From figure 2, you
can see that the example of two hydrogen atoms combing. As the atoms get closer
together, each electron experiences an attraction towards the two nuclei and the
electron density shifts so that the most probable place to find the two
electrons is between the two nuclei. Effectively each atom now has a share of
both the electrons. The electron density between the two nuclei exerts an
attractive force on each nucleus keeping them held tightly together in a
covalent bond.

Figure 2

A covalent bond forming between two hydrogen atoms.

It is also possible for two atoms share more than one pair of electrons, sharing
two pairs results in a double bond and sharing three pairs results in a triple
bond. Electronegativity is a measure of how powerful a atom is in a molecule to
attract electrons. Polarization is a term given to name the unequal sharing of
electrons in a covalent bond. Molecules that have unequal sharing of electrons
are called polar molecules and dipole molecules are ones which have the charge
separated, therefore all polar molecules must have a dipole attraction. Non-
polar molecules are ones in which there shapes are symmetrical so the electrons
are evenly distributed. Polar molecules have a permanent dipole in other words
they have a permanent separation of charge. As a result of this, polar molecules
are attracted to one another by forces called permanent dipole-permanent dipole
interactions, in which the negative end of one molecule is attracted towards the
positive end of another. These interactions decrease quite rapidly as the
distance between molecules increases. They are approximately 100 times weaker
than covalent bonds. There are also very strong types of dipole-dipole
interactions called Hydrogen bonds. Evidence for the existence of such
intermolecular forces lies in the properties of hydrides formed by element in
groups 4,5,6 and 7. While all the hydrides formed in group 5 behave in a similar
way, the hydrides of other groups do not. This suggest that the intermolecular
forces in these hydrides are much stronger than expected compared with other
hydrides of the other elements in each group. This type of intermolecular
bonding occurs in two molecules that each contain a polar bond between hydrogen
and another atom.

Figure 3

The variation in boiling points of the ...