Question

In each pair of compounds, pick the one with the higher boiling point. Explain your reasoning. MISSED THIS? Read Section 12.3; Watch \(\mathrm{KCV} 12.3,\) IWE 12.2 a. \(\mathrm{CH}_{3} \mathrm{OH}\) or \(\mathrm{CH}_{3} \mathrm{SH}\) b. \(\mathrm{CH}_{3} \mathrm{OCH}_{3}\) or \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\) c. \(\mathrm{CH}_{4}\) or \(\mathrm{CH}_{3} \mathrm{CH}_{3}\)

Short answer

a. CH3OH has a higher boiling point due to hydrogen bonding. b. CH3CH2OH has a higher boiling point due to hydrogen bonding. c. CH3CH3 has a higher boiling point due to stronger London dispersion forces.

Step-by-step solution

Understand Boiling Point Influences

Recognize that boiling points are influenced by molecular weight, structure, and the type and strength of intermolecular forces such as hydrogen bonding, dipole-dipole interactions, and London dispersion forces. The stronger the intermolecular forces, the higher the boiling point.

Analyze Compound a - CH3OH vs. CH3SH

Determine the dominant intermolecular force in each compound. Both methanol (CH3OH) and methanethiol (CH3SH) have similar molecular weights, but CH3OH can form hydrogen bonds due to the highly electronegative oxygen atom, whereas CH3SH has weaker dipole-dipole interactions from the less electronegative sulfur atom. Thus, CH3OH has the stronger intermolecular forces and a higher boiling point.

Analyze Compound b - CH3OCH3 vs. CH3CH2OH

Assess the intermolecular forces in the compounds. Dimethyl ether (CH3OCH3) is an ether with no capacity for hydrogen bonding, while ethanol (CH3CH2OH) can engage in hydrogen bonding due to its hydroxyl group. Even though both have similar molecular weights, ethanol's ability to hydrogen bond gives it a higher boiling point.

Analyze Compound c - CH4 vs. CH3CH3

Compare the molecular size and corresponding London dispersion forces. Methane (CH4) is smaller than ethane (CH3CH3). Although both are nonpolar and only London dispersion forces are present, the larger molecule, CH3CH3, has a greater surface area leading to stronger dispersion forces and hence a higher boiling point.

Key concepts

Intermolecular Forces

Intermolecular forces are the attractive forces that exist between individual molecules, influencing their physical state and properties such as boiling point. There are three primary types of intermolecular forces: hydrogen bonding, dipole-dipole interactions, and London dispersion forces.

Understanding these forces provides insight into why certain substances have higher boiling points. Generally, the stronger the intermolecular forces within a substance, the more energy is required in the form of heat to separate the molecules, resulting in a higher boiling point. Consider water (H₂O): its relatively high boiling point is due to strong hydrogen bonds between the water molecules, which are difficult to break.

Hydrogen Bonding

Hydrogen bonding is a strong type of dipole-dipole interaction, but it occurs specifically when a hydrogen atom is bonded to a highly electronegative atom, such as nitrogen, oxygen, or fluorine. The partially positive hydrogen is then attracted to a lone pair of electrons on a neighboring electronegative atom.

This type of bonding is particularly important in determining boiling points. For instance, in methanol (CH₃OH), the hydrogen-oxygen bond allows for the formation of hydrogen bonds, significantly raising its boiling point compared to compounds like methanethiol (CH₃SH), which cannot form such strong bonds. Due to hydrogen bonding, substances like water have unusually high boiling points compared to their molecular weight.

Dipole-Dipole Interactions

Dipole-dipole interactions occur between polar molecules, which have a permanent dipole moment due to a difference in electronegativity between the atoms involved. The partially positive end of one dipole is attracted to the partially negative end of a neighboring molecule.

These interactions are generally weaker than hydrogen bonds. An example that illustrates this is the comparison between dimethyl ether (CH₃OCH₃) and ethanol (CH₃CH₂OH). Both are polar, but only ethanol has the capacity for hydrogen bonding due to its hydroxyl group, resulting in a higher boiling point than dimethyl ether, which relies solely on dipole-dipole interactions.

London Dispersion Forces

London dispersion forces are the weakest intermolecular forces and are present in all molecules, polar or nonpolar. They arise from temporary dipoles that occur when the electron distribution within a molecule becomes asymmetric.

These forces increase with the size of the molecule, as larger molecules have more electrons and a larger surface area that can induce dipoles in adjacent molecules. Considering methane (CH₄) and ethane (CH₃CH₃), we see that ethane, being the larger molecule, possesses stronger London dispersion forces. Consequently, ethane has a higher boiling point than methane.