Question

In each of the following pairs, which would you expect to have the higher boiling point? (a) \(\mathrm{C}_{7} \mathrm{H}_{16}\) or \(\mathrm{C}_{10} \mathrm{H}_{22} ;\) (b) \(\mathrm{C}_{3} \mathrm{H}_{8}\) or \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{O} ;\) (c) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{SH}\) or \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\)

Short answer

The compounds with higher boiling points are: \(\mathrm{C}_{10} \mathrm{H}_{22}\), \((\mathrm{CH}_{3})_{2} \mathrm{O}\), and \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\).

Step-by-step solution

Analyze Pair (a)

Pair (a) includes two hydrocarbons: \(\mathrm{C}_{7} \mathrm{H}_{16}\) and \(\mathrm{C}_{10} \mathrm{H}_{22}\). In hydrocarbons, the primary intermolecular force is the London dispersion force. Since \(\mathrm{C}_{10} \mathrm{H}_{22}\) has more electrons and a larger surface area than \(\mathrm{C}_{7} \mathrm{H}_{16}\), this means \(\mathrm{C}_{10} \mathrm{H}_{22}\) has a higher boiling point.

Analyze Pair (b)

Pair (b) contains \(\mathrm{C}_{3} \mathrm{H}_{8}\) and \((\mathrm{CH}_{3})_{2} \mathrm{O}\). The first compound is a hydrocarbon, and the second is an ether. Though they both have London dispersion forces, the ether also has dipole-dipole interactions due to the presence of oxygen, which makes the forces between its molecules stronger than in \(\mathrm{C}_{3} \mathrm{H}_{8}\). Therefore, \((\mathrm{CH}_{3})_{2} \mathrm{O}\) has the higher boiling point.

Analyze Pair (c)

Pair (c) features \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{SH}\) and \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\). Even though both molecules have similar molar masses and both experience dipole-dipole interactions, the compound with OH group (\(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\)) also forms hydrogen bonds. This type of bonding is stronger than others, meaning \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\) has the higher boiling point.

Key concepts

Intermolecular Forces

When we talk about intermolecular forces, we're referring to the attractions that occur between molecules. They play a crucial role in determining the physical properties of substances, including boiling points. There are several types of these forces, but they can broadly be grouped into three categories: London dispersion forces, dipole-dipole interactions, and hydrogen bonds.

The strength of intermolecular forces depends on a variety of factors, such as the size of the molecules and the presence of polar bonds. The general rule is the stronger the intermolecular forces, the higher the boiling point. This is because more energy is needed to overcome the attraction between molecules in a liquid state to convert it into a gaseous state.

In a pair of simple hydrocarbons, like heptane (C7H16) and decane (C10H22), both exhibit London dispersion forces. The larger molecule, decane, with a greater number of electrons and surface area, demonstrates stronger dispersion forces and hence a higher boiling point compared to heptane.

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole interaction and is significantly stronger than other dipole interactions and London dispersion forces. This type of bonding occurs when hydrogen is covalently bonded to a highly electronegative atom such as fluorine, oxygen, or nitrogen, and this hydrogen is attracted to an electronegative atom in a neighboring molecule.

This attraction is responsible for the anomalously high boiling points of water (H2O), ammonia (NH3), and hydrogen fluoride (HF). An excellent example, taken from the exercise, is the comparison between ethanethiol (CH3CH2SH) and ethanol (CH3CH2OH). Both molecules have similar molar masses and exhibit dipole-dipole interactions due to their polar bonds. However, the OH group in ethanol can participate in hydrogen bonding, which increases the strength of the intermolecular forces present and, as a result, ethanol has the higher boiling point. The capacity for hydrogen bonding often makes substances with OH or NH groups much more likely to have higher boiling points than otherwise expected.

London Dispersion Forces

London dispersion forces, sometimes referred to as van der Waals forces, are the weakest type of intermolecular force but are universally present in all molecular substances. They arise due to the temporary polarization of electrons in a molecule, which induces a temporary dipole that can then induce dipoles in adjacent molecules.

The strength of London dispersion forces is influenced by the size of the molecules and the shape. Larger molecules with more electrons can become more polarizable and hence exhibit stronger dispersion forces. Similarly, elongated molecules that can have more surface contact with each other will have stronger dispersion forces compared to compact, spherical molecules.

In the exercise, basic hydrocarbons like propane (C3H8) and ethers like dimethyl ether ((CH3)2O) are compared. While both exhibit London dispersion forces, it’s important to note that the shape of the molecules and their ability to pack closely can influence the strength of these forces and, consequently, the boiling point. Despite the presence of dipole-dipole interactions in ethers, London dispersion forces play a significant role in determining their boiling point.