Question

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

Short answer

In short, the differences in boiling points for the given pairs can be explained as follows: a) HF has a higher boiling point than HCl due to stronger hydrogen bonding in HF compared to weaker dipole-dipole interactions in HCl. b) CHBr3 has a higher boiling point than CHCl3 due to larger bromine atoms, leading to stronger London dispersion forces. c) ICl has a higher boiling point than Br2 because ICl has additional dipole-dipole interactions in addition to London dispersion forces, resulting in stronger intermolecular forces than in Br2.

Step-by-step solution

Pair A: HF (20°C) and HCl (-85°C)

In Pair A, we have hydrogen fluoride (HF) with a boiling point of 20°C and hydrogen chloride (HCl) with a boiling point of -85°C. To explain the difference in boiling points, let's compare the intermolecular forces present in both molecules. HF forms strong hydrogen bonds due to the highly electronegative fluorine atom. These hydrogen bonds create strong intermolecular forces between HF molecules requiring more energy to break them apart, thus leading to a higher boiling point. In contrast, HCl molecules experience dipole-dipole interactions due to the difference in electronegativity between hydrogen and chlorine. However, the electronegativity difference is not as strong as it is for HF, and hence, these forces are weaker than the hydrogen bonds in HF. Consequently, HCl needs less energy to separate its molecules, thus resulting in a lower boiling point.

Pair B: CHCl3 (61°C) and CHBr3 (150°C)

In Pair B, we have chloroform (CHCl3) with a boiling point of 61°C and bromoform (CHBr3) with a boiling point of 150°C. To understand the difference in boiling points, let's look at the intermolecular forces and atomic size. Both CHCl3 and CHBr3 molecules have similar molecular structures and are polar in nature, which lead to dipole-dipole interactions between the molecules. However, CHBr3 has larger and heavier bromine atoms than the lighter chlorine atoms in CHCl3. As a result, the London dispersion forces (which increase with the molecular size) are stronger between the CHBr3 molecules than the CHCl3 molecules. This increased strength in London dispersion forces for CHBr3 requires more energy to overcome, leading to a higher boiling point than CHCl3.

Pair C: Br2 (59°C) and ICl (97°C)

In Pair C, we have bromine (Br2) with a boiling point of 59°C and iodine monochloride (ICl) with a boiling point of 97°C. To explain the difference in boiling points, let's compare the intermolecular forces and molecular structures. Bromine is a diatomic molecule (Br2) which experiences London dispersion forces due to its nonpolar nature. In contrast, iodine monochloride (ICl) is a polar molecule as there is a significant difference in electronegativity between iodine and chlorine atoms. Consequently, ICl experiences dipole-dipole interactions in addition to London dispersion forces. While both Br2 and ICl experience London dispersion forces, the additional dipole-dipole interactions in ICl leads to stronger intermolecular forces than in Br2. Thus, ICl has a higher boiling point than Br2 since more energy is required to overcome these stronger interactions between the ICl molecules.

Key concepts

Intermolecular Forces

When discussing the boiling point of a substance, one of the primary considerations is the nature and strength of intermolecular forces. Intermolecular forces are the attractive forces between molecules, and they play a critical role in determining the physical properties of substances, including their boiling points.

There are three main types of intermolecular forces:

• Hydrogen bonding, a special type of strong dipole-dipole attraction that occurs when hydrogen is bonded to highly electronegative elements like oxygen, nitrogen, or fluorine.
• Dipole-dipole interactions, which occur between polar molecules where there is an uneven distribution of electrical charge.
• London dispersion forces, which are the weakest and arise due to momentary changes in electron density in nonpolar molecules or atoms.

Substances with stronger intermolecular forces require more energy to separate the molecules, which translates to a higher boiling point.

Hydrogen Bonding

Hydrogen bonding is one of the strongest forms of intermolecular attraction and has a profound impact on the boiling points of compounds. It occurs when a hydrogen atom is covalently bonded to a highly electronegative atom like fluorine, oxygen, or nitrogen, and is also electrostatically attracted to the lone pair of another electronegative atom in a neighboring molecule.

This type of bonding can be likened to a powerful magnet that holds molecules together tightly, requiring substantial energy to break. For example, in the case of HF, the presence of hydrogen bonding explains its significantly higher boiling point compared to HCl, which does not form hydrogen bonds and relies on the weaker dipole-dipole interactions due to its polarity.

London Dispersion Forces

London dispersion forces are the weakest type of intermolecular force and are present in all molecules, whether they are polar or nonpolar. They arise due to the momentary uneven distribution of electrons around a molecule, leading to temporary dipoles that attract each other. The strength of London dispersion forces increases with larger atoms or molecules, as there is a greater area over which the electron cloud can become distorted.

In the boiling point comparison between CHCl₃ and CHBr₃, the larger size and mass of bromine compared to chlorine result in stronger London dispersion forces in CHBr₃, giving it a higher boiling point despite both molecules being polar and exhibiting dipole-dipole interactions as well.

Molecular Polarity

Molecular polarity refers to the distribution of electrical charge over the atoms in a molecule. A molecule is polar if there is a significant difference in the electronegativity between the atoms, resulting in a partial positive charge on one end and a partial negative charge on the other end. This uneven charge distribution leads to dipole-dipole interactions, where the positive end of one molecule attracts the negative end of another.

In the case of ICl, its boiling point is higher than that of Br₂ because ICl is polar, leading to stronger dipole-dipole forces in addition to London dispersion forces. This combined effect of molecular polarity increases the intermolecular forces and, as a result, the boiling point.

Electronegativity Difference

The difference in electronegativity between atoms in a molecule can determine whether a molecule will be polar or nonpolar, which in turn influences intermolecular forces and boiling points. Electronegativity is the ability of an atom to attract electrons to itself. When two atoms bonded together have a significant electronegativity difference, the bond becomes polar, often resulting in a polar molecule.

For instance, the HF molecule has a high electronegativity difference between hydrogen and fluorine, leading to a strong polarity and the capability to form hydrogen bonds, which are stronger than other dipole-dipole interactions. This difference in electronegativity is key to understanding why HF has a higher boiling point compared to molecules with a smaller electronegativity difference, such as HCl.