Identify the intermolecular forces present in the following solids: (a) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\) (b) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{3}\) (c) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{Cl}\)

When discussing intermolecular forces, hydrogen bonding often emerges as one of the strongest kinds, second only to ionic and covalent bonds in terms of strength. It occurs when a hydrogen atom, covalently bonded to a highly electronegative atom such as oxygen, fluorine, or nitrogen, experiences an attraction towards a nearby electronegative atom on another molecule. This special interaction is what gives water its remarkable properties, like its high boiling point relative to its molecular weight.

Let's take ethanol (\texttt{CH\(_{3}\)CH\(_{2}\)OH}) as an example, which was mentioned in the exercise. Its molecular structure includes an -OH group, where the highly electronegative oxygen forms a dipolar bond with hydrogen. This results in partial positive and negative charges, enabling the formation of hydrogen bonds with other polar molecules. In solids, these forces create a network of connections, heavily influencing physical properties such as melting and boiling points.

In educational materials, hydrogen bonding is often represented by dotted lines connecting hydrogen to the electronegative atoms of adjacent molecules to illustrate this unique interaction that goes beyond ordinary dipole-dipole interactions.

Dipole-dipole interactions are another type of intermolecular force where the positive end of a polar molecule is attracted to the negative end of another polar molecule. These attractions occur between polar molecules that have permanent dipole moments. For instance, when we look at chloroethane (\texttt{CH\(_{3}\)CH\(_{2}\)Cl}), it has a polar C-Cl bond due to the significant difference in electronegativity between carbon and chlorine.

These permanent dipoles align such that oppositely charged ends are close to each other. The strength of dipole-dipole interactions is greatly dependent on the polar nature of the molecule and the physical proximity of the molecules. While they are weaker than hydrogen bonds, they play a crucial role in determining the physical properties of substances such as their phase at room temperature and solubility in different solvents.

London dispersion forces, also known as Van der Waals forces, are the weakest but most ubiquitous type of intermolecular force. They are present in all molecular substances, whether polar or nonpolar. These forces originate from temporary dipoles that are created when the electrons in a molecule are unevenly distributed at a particular instant in time, leading to a temporary separation of charge.

These temporary dipoles induce others in neighboring molecules, creating an instantaneous attraction. Even in nonpolar molecules like propane (\texttt{CH\(_{3}\)CH\(_{2}\)CH\(_{3}\)}), which lacks polar bonds and hydrogen bonding, London dispersion forces are present. These forces are more prominent in larger molecules with more electrons, and they can contribute to various physical properties such as boiling points and viscosities. Despite their relatively weak nature, London dispersion forces are significant because of their universal presence in molecules.