Question

Rationalize the difference in boiling points for each of the following pairs of substances: $$\begin{array}{rr}{\text { a. Ar }} & {-186^{\circ} \mathrm{C}} \\\ {\mathrm{HCl}} & {-85^{\circ} \mathrm{C}}\end{array}$$ $$\begin{array}{rr}{\text { b. } \mathrm{HF}} & {20^{\circ} \mathrm{C}} \\\ {\mathrm{HCl}} & {-85^{\circ} \mathrm{C}}\end{array}$$ $$\begin{array}{cc}{\text { c. } \mathrm{HCl}} & {-85^{\circ} \mathrm{C}} \\\ {\mathrm{LiCl}} & {1360^{\circ} \mathrm{C}}\end{array}$$ $$\begin{array}{ccc}{\text { d. } n \text { -pentane }} & {\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{3}} & {36.2^{\circ} \mathrm{C}} \\ {n \text { -hexane }} & {\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{3}} & {69^{\circ} \mathrm{C}}\end{array}$$

Short answer

In summary, the differences in boiling points for the given pairs of substances can be explained by the type and strength of the intermolecular forces present in these molecules: a. Ar has weak London dispersion forces, while HCl has stronger dipole-dipole interactions and hydrogen bonding, leading to a higher boiling point for HCl. b. Both HF and HCl exhibit dipole-dipole interactions, but the hydrogen bonding in HF is much stronger due to fluorine's higher electronegativity, resulting in a higher boiling point for HF. c. HCl has dipole-dipole interactions and hydrogen bonding, while LiCl has strong ionic bonds, which require much more energy to break, causing a higher boiling point for LiCl. d. n-pentane and n-hexane have London dispersion forces, but n-hexane's longer carbon chain and higher molecular weight result in stronger forces and a higher boiling point compared to n-pentane.

Step-by-step solution

Pair a: Ar vs HCl

: We have two substances Ar and HCl boiling at -186°C and -85°C, respectively. Ar is a noble gas, which means it has a complete valence electron shell and no tendency to react with other elements. The only intermolecular forces present in Ar are weak London dispersion forces (also known as van der Waals forces). HCl is a polar covalent molecule due to the electronegativity difference between H and Cl atoms. In hydrogen chloride molecules, there are strong dipole-dipole interactions and hydrogen bonding in addition to London dispersion forces. Overall, the intermolecular forces in HCl are stronger than in Ar. The difference in boiling points can be rationalized by the fact that stronger intermolecular forces require more energy to break, leading to a higher boiling point for HCl compared to Ar.

Pair b: HF vs HCl

: We have two substances HF and HCl boiling at 20°C and -85°C, respectively. Both HF and HCl are polar covalent molecules due to the electronegativity difference between H and the halogens (F and Cl). Both molecules exhibit dipole-dipole interactions; however, the hydrogen bonding in HF is much stronger than in HCl because fluorine is more electronegative than chlorine. The stronger hydrogen bonding in HF results in higher boiling point as it requires more energy to break these strong intermolecular forces.

Pair c: HCl vs LiCl

: We have two substances HCl and LiCl boiling at -85°C and 1360°C, respectively. As mentioned earlier, HCl is a polar covalent molecule with dipole-dipole interactions and hydrogen bonding in addition to London dispersion forces. LiCl is an ionic compound, which means it's composed of positively charged lithium ions (Li+) and negatively charged chloride ions (Cl-). The intermolecular forces present in LiCl are ionic bonds, which are much stronger than the forces in HCl. The difference in boiling points can be rationalized by the fact that breaking ionic bonds in LiCl requires much more energy than overcoming the intermolecular forces present in HCl.

Pair d: n-pentane vs n-hexane

: We have two substances n-pentane and n-hexane boiling at 36.2°C and 69°C, respectively. n-pentane and n-hexane are both non-polar hydrocarbons due to the small electronegativity difference between C and H atoms. The only intermolecular forces present in these molecules are London dispersion forces. The difference in boiling points can be explained by the difference in molecule size and molecular weight. n-hexane has a longer carbon chain and a higher molecular weight than n-pentane, which means it has a larger surface area and stronger London dispersion forces. These stronger intermolecular forces in n-hexane require more energy to break, resulting in a higher boiling point compared to n-pentane.

Key concepts

Intermolecular Forces

Intermolecular forces are the attractive forces that act between molecules. They are crucial for determining the physical properties of substances, such as boiling points. These forces vary in strength and type:

• London Dispersion Forces: Present in all molecules, but dominate in non-polar molecules like Argon (Ar) and hydrocarbons.
• Dipole-Dipole Interactions: Occur in polar molecules, such as hydrogen chloride (HCl), where there is a difference in electronegativity between atoms.
• Hydrogen Bonding: A special type of dipole-dipole interaction found in molecules with a hydrogen atom covalently bonded to a highly electronegative atom like fluorine, oxygen, or nitrogen, exemplified in HF.
• Ionic Bonds: Strong forces found in ionic compounds like lithium chloride (LiCl), which consist of charged ions.

These intermolecular forces require energy to break in order for a substance to transition from a liquid to a gas, affecting its boiling point directly. Stronger forces lead to higher boiling points.

Hydrogen Bonding

Hydrogen bonding is a powerful type of dipole-dipole interaction that has a significant impact on boiling points. It occurs when a hydrogen atom, which is bonded to a highly electronegative atom such as fluorine, experiences an attractive force from an electronegative atom in a nearby molecule. This bonding is seen in substances like hydrogen fluoride (HF).

• Strength: The strength of hydrogen bonds is greater than other dipole-dipole interactions due to the small size of hydrogen, which allows for close approach to electronegative atoms.
• Boiling Point Effect: Hydrogen bonding is responsible for HF having a higher boiling point than HCl, even though both are polar molecules. Fluorine's greater electronegativity compared to chlorine leads to stronger hydrogen bonds.

The presence of hydrogen bonds in a substance usually implies that significant energy is required to separate the molecules, resulting in higher boiling points.

Ionic Bonds

Ionic bonds are among the strongest types of intermolecular forces, formed between positively and negatively charged ions. These bonds are evident in ionic compounds like lithium chloride (LiCl).

• Composition: Ionic compounds consist of a metal and a non-metal, where electrons are transferred from the metal to the non-metal, creating positively charged cations and negatively charged anions.
• Properties: Due to the strong electrostatic attraction between ions, ionic compounds generally have high boiling and melting points.
• Comparison with Covalent Bonds: Compared to covalent compounds like hydrogen chloride (HCl), ionic compounds such as LiCl require much more energy to change state from solid to liquid or liquid to gas. This results in the markedly high boiling point of LiCl.

Thus, the robust nature of ionic bonds significantly elevates the boiling points of compounds in which they are present.

London Dispersion Forces

London dispersion forces are the weakest intermolecular forces but are universally present, especially in non-polar molecules. They arise from temporary fluctuations in electron density, which induce instantaneous dipoles.

• Occurrence: All molecules experience these forces, but they are the sole intermolecular force in noble gases like Argon (Ar) and non-polar compounds, such as n-pentane and n-hexane.
• Influence on Boiling Points: The strength of London dispersion forces is correlated with the size of the molecules and the surface area. Larger molecules like n-hexane have stronger dispersion forces compared to smaller ones like n-pentane, leading to higher boiling points.
• Effect of Structure: The more elongated the molecule, the greater the surface area and the stronger the dispersion forces, contributing to a higher boiling point.

In summary, while London dispersion forces are the weakest, they can still significantly impact boiling points in larger non-polar molecules.