Question

Ionic compounds tend to have higher boiling points than covalent substances do. Both ammonia, \(\mathrm{NH}_{3}\), and methane, \(\mathrm{CH}_{4}\), are covalent compounds, yet the boiling point of ammonia is \(130^{\circ} \mathrm{C}\) higher than that of methane. What might account for this large difference?

Short answer

The large difference in boiling points is due to hydrogen bonding in ammonia, which is much stronger than the dispersion forces in methane.

Step-by-step solution

- Understand the Types of Compounds

Identify that both ammonia \(\text{NH}_{3}\) and methane \(\text{CH}_{4}\) are covalent compounds. These compounds share electrons rather than transferring them like ionic compounds.

- Compare Boiling Points

Observe that ammonia has a boiling point that is \(\text{130}^{\text{\circ}}\text{C}\) higher than methane. This indicates a stronger intermolecular force in ammonia compared to methane.

- Intermolecular Forces in Ammonia

Ammonia molecules exhibit hydrogen bonding due to the presence of hydrogen atoms bonded to a highly electronegative nitrogen atom. Hydrogen bonds are strong intermolecular forces which lead to higher boiling points.

- Intermolecular Forces in Methane

Methane, on the other hand, primarily exhibits weak London dispersion forces because it is a nonpolar molecule with no significant electron density differences. Thus, methane has a lower boiling point.

- Conclusion

Summarize that the large difference in boiling points between ammonia and methane is due to the presence of hydrogen bonding in ammonia, which is much stronger than the dispersion forces present in methane.

Key concepts

Covalent Compounds

Covalent compounds are chemical substances formed by atoms sharing electrons. This occurs because these atoms have similar electronegativities and do not easily give up or accept electrons to form ionic bonds. The molecules in covalent compounds are held together by covalent bonds, which are strong and stable.
Some common properties of covalent compounds include:

• Low melting and boiling points compared to ionic compounds.
• They usually do not conduct electricity in any state.
• They can be gases, liquids, or solids at room temperature.

Examples include water \(\text{H}_{2}\text{O}\), carbon dioxide \(\text{CO}_{2}\), and methane \(\text{CH}_{4}\). While they share electrons, the intermolecular forces between their molecules influence their physical properties.

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like nitrogen, oxygen, or fluorine. This creates a strong attraction between the hydrogen atom of one molecule and the electronegative atom of another molecule.
For example, in ammonia \(\text{NH}_{3}\), the nitrogen atom is highly electronegative and forms a strong hydrogen bond with a hydrogen atom of another ammonia molecule. This strong interaction significantly influences physical properties like boiling points.
Here are some critical points to remember about hydrogen bonding:

• It is stronger than Van der Waals forces but weaker than covalent or ionic bonds.
• It leads to higher boiling and melting points in compounds that exhibit hydrogen bonding.
• It plays a crucial role in the structure and function of biological molecules like DNA and proteins.

London Dispersion Forces

London dispersion forces, also known as induced dipole-induced dipole forces, are the weakest type of intermolecular force. They arise from temporary fluctuations in electron density within molecules, leading to a momentary dipole.
These forces are present in all molecules, whether polar or nonpolar, but are the dominant force in nonpolar molecules like methane \(\text{CH}_{4}\). Because these forces are weak, substances that rely on London dispersion forces usually have lower boiling and melting points.
Key characteristics of London dispersion forces include:

• They increase with the number of electrons in a molecule.
• They are significant in larger atoms and molecules due to the larger electron cloud.
• They become stronger at colder temperatures due to reduced kinetic energy of molecules.

Understanding these forces helps explain why nonpolar substances like methane have much lower boiling points compared to substances such as ammonia, where stronger intermolecular forces are at play.