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Intermolecular forces

Van der Waals Interactions
As described earlier, there exist intermolecular attractions in the liquid state of the substance. The nature of these factors of attraction varies from molecule to molecule. These factors are much weaker than the valence forces which are responsible for uniting atoms via the formation of covalent or ionic bonds. The forces which keep the neutral molecules together in the liquid or solid phases are collectively known as Van der Waals forces (or bonds).

As mentioned above, the Van der Waals bonding is much weaker than covalent bonding. For example, the enthalpy of vaporization of methane (which measures the strength of the intermolecular bonds in the liquid) is only 8.2 kJ mol-1 as compared to the value of 435 kJ mol-1 for the bond dissociation energy of the C-H bond. Moreover, the Van der Waals bond operates over a longer distance as compared to covalent bonds. For example, in solid iodine, the molecules are at distance of 430 pm from each other as compared to a value of 266 pm for the interatomic distance. Broadly, Van der Waals forces may be classified into three categories, described as follows.

Dipole-Dipole Forces (Keesom, 1912)
If the covalent molecule has a permanent dipole moment, then the positively-charged end of the dipole of one molecule will attract the negatively-charged end of another molecule and thus molecules will have specific orientation with respect to each other. The contribution made to the net Van der Waals attraction by the dipole-dipole force is usually small.

Dipole-Induced Dipole Forces (Debye, 1920)
A molecule with a permanent dipole moment can induce a dipole in another molecule and this is followed by an attraction between the two molecules. The contribution made to the van der Waals interaction from this effect is also very small.

Dispersion Forces (London, 1930)
The above two forces cannot explain the forces of attraction which exist between the molecules of inert gases (helium, neon and argon). The forces of attraction in such molecules (and also in other types of molecules) are explained on the basis that the motion of an electronic cloud at any instant may be slightly displaced relative to the positive nucleus. The atom thus acquires an instantaneous dipole moment which can induce another atom nearby and, hence, an instantaneous net attraction exists between the two atoms (fig). 

However, the instantaneous dipole will continuously change its charge orientation so that over a definite interval of time (which is large compared to the instantaneous time) the average dipole moment of the molecule is zero. The instantaneous dipole-instantaneous induced dipole interaction forces are known as dispersion or London forces. This force of attraction occurs for only an extremely small time interval and also it acts over a very short distance. For this reason, these forces are called short-range forces.

The dispersion forces depend on the size of the molecule. In larger molecules, the electronic cloud extends over a larger space and, hence, they can also undergo a larger momentary distortion. Thus dispersion forces between molecules increase in the relative molecular mass.

Of the three forces of attraction, the dispersion forces contribute a major share towards the net Van der Waals interactions between molecules.

Theoretical calculations have shown that each of the three interactions described above varies inversely as the sixth power of the distance between the molecules. For a very small distance where electron clouds of two molecules begin to overlap, the repulsive forces start operating. These repulsive forces vary inversely as the twelfth power of the intermolecular distance. The net potential energy of interaction is given by the expression

where a and b are constants of proportionality. The above relationship with its components is shown in fig.


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