Question

Rationalize the difference in boiling points between the members of the following pairs of substances: (a) HF \(\left(20^{\circ} \mathrm{C}\right)\) and \(\mathrm{HCl}\left(-85^{\circ} \mathrm{C}\right)\), (b) \(\mathrm{CHCl}_{3}\left(61{ }^{\circ} \mathrm{C}\right)\) and \(\mathrm{CHBr}_{3}\) \(\left(150^{\circ} \mathrm{C}\right)\), (c) \(\mathrm{Br}_{2}\left(59^{\circ} \mathrm{C}\right)\) and \(\mathrm{ICl}\left(97^{\circ} \mathrm{C}\right)\).

Short answer

In summary, the differences in boiling points for the three pairs of substances are due to the types and strengths of intermolecular forces experienced by each substance. For pair (a), HF has stronger hydrogen bonding and dipole-dipole forces compared to HCl, leading to a higher boiling point. In pair (b), CHBr3 experiences stronger van der Waals forces compared to CHCl3 due to its larger molar mass, resulting in a higher boiling point. For pair (c), ICl has stronger dipole-dipole forces in addition to London dispersion forces compared to Br2, which only experiences London dispersion forces, leading to a higher boiling point for ICl.

Step-by-step solution

Pair (a): HF and HCl

To rationalize the difference in boiling points between HF and HCl, we need to compare their intermolecular forces. Both of these substances experience dipole-dipole forces due to the difference in electronegativity between their atoms. However, HF also has hydrogen bonding, which is a stronger intermolecular force than dipole-dipole forces. The electronegativity difference between H and F is higher compared to H and Cl, which causes stronger dipole-dipole forces between HF molecules. Additionally, hydrogen bonding in HF occurs because F is one of the most electronegative elements and can form hydrogen bonds with H atoms. The stronger intermolecular forces in HF require more energy to overcome, leading to its higher boiling point compared to HCl.

Pair (b): CHCl3 and CHBr3

In order to compare the boiling points of CHCl3 and CHBr3, let's consider the types of intermolecular forces present in both substances. Both of these molecules are polar and experience dipole-dipole forces. However, the difference in boiling points is mainly due to the difference in van der Waals forces (or London dispersion forces) between the molecules. The molar masses of CHBr3 and CHCl3 are 271 g/mol and 119.38 g/mol, respectively. The larger molar mass of CHBr3 implies a greater number of electrons, which in turn results in stronger van der Waals forces. The stronger van der Waals forces in CHBr3 require more energy to overcome, leading to its higher boiling point compared to CHCl3.

Pair (c): Br2 and ICl

To explain the difference in boiling points between Br2 and ICl, let's analyze their intermolecular forces. Both Br₂ and ICl are nonpolar molecules, and as such, they have London dispersion forces (or van der Waals forces) between their molecules. In the case of Br₂, it is a homonuclear diatomic molecule with a molar mass of 159.8 g/mol, while ICl is a heteronuclear diatomic molecule with a molar mass of 162.36 g/mol. Although their molar masses are very similar, the primary difference lies in their molecular polarity. ICl is a polar molecule due to the electronegativity difference between I and Cl, leading to stronger dipole-dipole forces in addition to the London dispersion forces. In the case of Br2, only London dispersion forces are present as it is a nonpolar molecule. The presence of stronger dipole-dipole forces in ICl requires more energy to overcome, resulting in its higher boiling point compared to Br2.

Key concepts

Boiling Points Comparison

Understanding the boiling points of different substances is crucial in the realm of chemistry. In general, the boiling point of a substance is the temperature at which the vapor pressure equals the atmospheric pressure, allowing the substance to change from a liquid to a gas. Comparing boiling points reveals a lot about a substance’s intermolecular forces — the attractive forces between neighboring molecules.

Stronger intermolecular forces result in higher boiling points because more energy is needed to separate molecules from each other. For example, in the comparison between HF and HCl, the significant difference in boiling points can be attributed to hydrogen bonding in HF, which is a much stronger force than the dipole-dipole interactions in HCl. Similarly, variations in intermolecular forces explain the boiling point differences between CHCl3 and CHBr3, and between Br2 and ICl.

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole interaction that occurs when hydrogen is bonded to a highly electronegative atom like nitrogen, oxygen, or fluorine. These bonds are incredibly strong compared to regular dipole-dipole interactions because the electronegative atom pulls the electron density away from hydrogen, leaving it with a significant partial positive charge. This charge then strongly attracts the electronegative atom in neighboring molecules.

In the case of HF, the fluorine atom’s high electronegativity facilitates the formation of hydrogen bonds, which lead to a much higher boiling point compared to HCl, where such bonding is absent. This is why HF has a remarkably high boiling point for such a small molecule. Incorporating examples like water and ammonia can further reinforce students’ understanding of hydrogen bonding's significant impact on physical properties.

Van der Waals Forces

Van der Waals forces encompass a variety of weak intermolecular forces, including dipole-dipole interactions and London dispersion forces. The latter are temporary dipoles created when electron distribution within a molecule is uneven and induces a dipole in a neighboring molecule. This can happen in any molecule, whether polar or nonpolar, but it's especially significant in larger, heavier atoms and molecules.

The boiling points of CHCl3 and CHBr3, as well as Br2 and ICl, are influenced by these forces. The higher molar mass and the resultant increase in electrons of CHBr3, compared to CHCl3, lead to stronger London dispersion forces, and thus, a higher boiling point. These forces are also the reason for the difference in boiling points between the nonpolar Br2 and the polar ICl, with the latter also supporting dipole-dipole interaction due to its molecular polarity.

Molecular Polarity

Molecular polarity is determined by the distribution of electrical charge over the atoms in a molecule. For a molecule to be polar, it must have polar bonds (bonds with a difference in electronegativity between the bonded atoms), and the shape of the molecule should allow for an uneven distribution of charge. Polarity significantly influences the physical properties of a substance, such as boiling point and melting point.

In our example, ICl is polar due to the difference in electronegativity between iodine and chlorine, thus experiencing both London dispersion and dipole-dipole forces, leading to a higher boiling point compared to the nonpolar Br2. Similarly, when students understand the concept of molecular polarity, they can better grasp why molecules with polar bonds like HF exhibit much stronger intermolecular forces and higher boiling points relative to nonpolar molecules such as HCl.