Question

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 \(\mathrm{ICl}\left(97^{\circ} \mathrm{C}\right)\)

Short answer

The difference in boiling points for each pair can be explained by the intermolecular forces present: (a) HF has a higher boiling point due to stronger hydrogen bonding compared to HCl's dipole-dipole forces, (b) CHBr3 has a higher boiling point due to stronger London dispersion forces and larger dipole moment compared to CHCl3, and (c) ICl has a higher boiling point due to the presence of both dipole-dipole forces and London dispersion forces compared to Br2's London dispersion forces alone.

Step-by-step solution

Understanding Intermolecular Forces

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.

Pair (a): HF and HCl

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).

Pair (b): CHCl3 and CHBr3

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).

Pair (c): Br2 and ICl

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.

Key concepts

Dipole-Dipole Forces

Dipole-dipole forces are a type of intermolecular force that occur between polar molecules. These forces exist because of the electrostatic attraction between the positive and negative ends of polar molecules.
When two molecules have permanent dipoles, they align so that the positive end of one molecule is close to the negative end of another.
This alignment creates a force of attraction, known as dipole-dipole forces, which can significantly affect the substance's properties.

• Strength: Dipole-dipole forces are stronger than London dispersion forces but weaker than hydrogen bonds.
• Impact: Substances with strong dipole-dipole interactions tend to have higher boiling points compared to those with only London dispersion forces.

Molecules like HCl, with a permanent dipole due to its polar nature, experience these forces, contributing to its physical characteristics, such as its boiling point.

Hydrogen Bonding

Hydrogen bonding is a special, highly potent form of dipole-dipole attraction. It specifically occurs when a hydrogen atom is covalently bonded to one of the three highly electronegative atoms – fluorine, oxygen, or nitrogen.
This creates a strong attraction because the electronegative atoms pull the electron density towards themselves, making the hydrogen atom very positive relative to the other end.

• Examples: Water (\( ext{H}_2 ext{O}\)), ammonia (\( ext{NH}_3\)), and hydrogen fluoride (\( ext{HF}\)) are classic examples where hydrogen bonding is prevalent.
• Effect: Hydrogen bonding greatly increases the boiling and melting points of substances since more energy is required to break these strong interactions.

In the pair HF and HCl, HF exhibits hydrogen bonding due to the hydrogen-fluorine bond, resulting in a much higher boiling point than HCl, which lacks this type of interaction.

London Dispersion Forces

London dispersion forces are the weakest type of intermolecular force. They are temporary forces that arise from momentary fluctuations in electron density within molecules.
These fluctuations create temporary dipoles, which induce similar dipoles in adjacent molecules, resulting in a weak attraction.

• Universality: These forces occur in all molecules, whether they are polar or nonpolar.
• Size and Shape Dependency: The strength of London dispersion forces increases with the size and surface area of the molecule. Larger atoms and more elongated shapes tend to have stronger dispersion forces.

In comparing CHCl \(_3\) with CHBr \(_3\) and Br \(_2\) with ICl, larger molecules like CHBr \(_3\) and ICl exhibit stronger London dispersion forces due to their increased size and larger electron clouds, contributing to their higher boiling points.