Rationalize the difference in boiling points in each pair: (a) HF \(\left(20^{\circ} \mathrm{C}\right)\) and \(\mathrm{HCl}\left(-85^{\circ} \mathrm{C}\right),\) (b) CHCl \(_{3}\left(61^{\circ} \mathrm{C}\right)\) and \(\mathrm{CHBr}_{3}\left(150^{\circ} \mathrm{C}\right),(\mathbf{c}) \mathrm{Br}_{2}\left(59^{\circ} \mathrm{C}\right)\) and \(\mathrm{ICl}\left(97^{\circ} \mathrm{C}\right)\)

Intermolecular forces are the forces between molecules in a substance. These forces are responsible for the various physical properties of substances, including boiling points. There are three main types of intermolecular forces: dipole-dipole forces, hydrogen bonding, and London dispersion forces. - Dipole-dipole forces occur between polar molecules with positive and negative ends. The stronger the dipole, the stronger the force. - Hydrogen bonding is a special case of dipole-dipole interaction that occurs between molecules containing a hydrogen atom bonded to a highly electronegative atom (usually fluorine, oxygen, or nitrogen). Hydrogen bonding is typically stronger than other dipole-dipole forces. - London dispersion forces are temporary forces that occur between all molecules, regardless of their polarity. These forces arise from temporary fluctuations in electron density and increase with the size and shape of the molecule. With this understanding, let's analyze each pair of compounds.

Both HF and HCl are diatomic molecules and have similar structures. However, HF exhibits hydrogen bonding due to the presence of a highly electronegative fluorine atom, while HCl does not. Hydrogen bonding is a stronger force than dipole-dipole forces present in HCl, leading to a higher boiling point for HF (20°C) as compared to HCl (-85°C).

Both of these molecules are similar in structure, but the presence of bromine atoms in CHBr3 makes it larger and heavier compared to CHCl3. This results in stronger London dispersion forces for CHBr3 compared to CHCl3. In addition, the dipole moment of CHBr3 is larger than that of CHCl3 due to the increased size and polarizability of the bromine atoms. These stronger intermolecular forces contribute to the higher boiling point of CHBr3 (150°C) compared to CHCl3 (61°C).

Br2 is a diatomic molecule with only London dispersion forces, while ICl is a polar molecule with both dipole-dipole forces and London dispersion forces present. The dipole-dipole forces present in ICl make its intermolecular forces stronger overall than those in Br2. This leads to a higher boiling point for ICl (97°C) compared to Br2 (59°C). In conclusion, the differences in boiling points for each pair of compounds can be rationalized by analyzing the intermolecular forces present in the molecules. Stronger intermolecular forces result in higher boiling points, as these forces must be overcome for the substance to change from a liquid to a gaseous state.