Question

In each pair of compounds, pick the one with the higher boiling point. Explain your reasoning. a. \(\mathrm{NH}_{3}\) or \(\mathrm{CH}_{4}\) b. \(\mathrm{CS}_{2}\) or \(\mathrm{CO}_{2}\) c. \(\mathrm{CO}_{2}\) or \(\mathrm{NO}_{2}\)

Short answer

NH3 has a higher boiling point than CH4 due to hydrogen bonding. CS2 has a higher boiling point than CO2 due to its greater molar mass and the resulting stronger London dispersion forces. NO2 has a higher boiling point than CO2 because it has dipole-dipole interactions in addition to London dispersion forces.

Step-by-step solution

Analyze Intermolecular Forces for NH3 and CH4

Examine the types of intermolecular forces (IMFs) present in ammonia (NH3) and methane (CH4). Ammonia molecules have hydrogen bonds, which are strong IMFs due to the presence of a highly electronegative nitrogen atom bonded to hydrogen. Methane has only weak London dispersion forces because it is a nonpolar molecule.

Compare Boiling Points for NH3 and CH4

Since stronger IMFs give rise to higher boiling points and ammonia has hydrogen bonding while methane does not, NH3 will have the higher boiling point than CH4.

Analyze Intermolecular Forces for CS2 and CO2

Identify the IMFs in carbon disulfide (CS2) and carbon dioxide (CO2). Both are nonpolar molecules and thus only have London dispersion forces. However, the larger molar mass of CS2 indicates stronger London dispersion forces and hence a higher boiling point.

Compare Boiling Points for CS2 and CO2

Since CS2 has stronger dispersion forces as compared to CO2 due to a greater molar mass, CS2 will have a higher boiling point.

Analyze Intermolecular Forces for CO2 and NO2

Determine the IMFs in CO2 and NO2. CO2 is nonpolar with only London dispersion forces, while NO2 is a polar molecule with dipole-dipole interactions in addition to London dispersion forces.

Compare Boiling Points for CO2 and NO2

Given that NO2 has stronger IMFs due to its polarity, which includes both London dispersion forces and dipole-dipole interactions, NO2 will have a higher boiling point than CO2.

Key concepts

Hydrogen Bonding

Understanding hydrogen bonding is crucial when it comes to predicting boiling points of molecules. This type of intermolecular force occurs when a hydrogen atom, covalently bonded to a highly electronegative atom such as nitrogen, oxygen, or fluorine, experiences attraction towards another electronegative atom with a lone pair of electrons.

In the context of our exercise, ammonia (\texttt{NH}\(_3\)) is an excellent example of a molecule that exhibits hydrogen bonding due to the presence of a nitrogen atom bonded to hydrogen. This gives it a significantly higher boiling point compared to methane (\texttt{CH}\(_4\)), which lacks this strong force and relies only on weaker London dispersion forces.

In general, the stronger the hydrogen bonding, the higher the boiling point, since more energy is required to break these interactions during the phase transition from liquid to gas.

London Dispersion Forces

London dispersion forces are the weakest type of intermolecular force, but they are nonetheless present in all molecules. They result from temporary fluctuations in electron distribution, which create temporary dipoles that induce attractions between adjacent molecules.

These forces are prominent in nonpolar molecules and increase with greater molar mass, as we see with carbon disulfide (\texttt{CS}\(_2\)) compared to carbon dioxide (\texttt{CO}\(_2\)). Even though both molecules only exhibit London dispersion forces, \texttt{CS}\(_2\) has a higher molar mass, leading to stronger intermolecular attractions and, therefore, a higher boiling point.

Essentially, heavier molecules or those with larger surface area tend to have higher boiling points due to stronger London dispersion forces.

Molecular Polarity

Molecular polarity refers to the distribution of electrical charge over the atoms in a molecule. A molecule is considered polar if there is a significant difference in electronegativity between its atoms, leading to an uneven distribution of electron density and thus a permanent dipole moment.

When comparing \texttt{CO}\(_2\) and \texttt{NO}\(_2\), the latter is polar because of the uneven charge distribution and the bent molecular geometry. As a result, \texttt{NO}\(_2\) molecules are attracted to each other by dipole-dipole interactions, which are absent in nonpolar \texttt{CO}\(_2\). This polarity and the resulting intermolecular forces contribute to \texttt{NO}\(_2\)'s higher boiling point as it requires more energy to overcome these forces during boiling.

Dipole-Dipole Interactions

Dipole-dipole interactions are intermolecular forces that occur between polar molecules. These interactions are due to the attraction between the positive end of one polar molecule and the negative end of another. The strength of dipole-dipole interactions is affected by the polarity of the molecules involved; the more polar the molecule, the stronger the interaction.

In our exercise, we consider \texttt{NO}\(_2\), a molecule with a significant dipole moment due to its polar nature, leading to dipole-dipole interactions in addition to London dispersion forces. This makes the boiling point of \texttt{NO}\(_2\) higher than that of \texttt{CO}\(_2\). The presence of these dipole-dipole interactions is a key factor in determining a molecule's boiling point and must be considered alongside other types of intermolecular forces.