A The boiling points of HCl, HBr, and HI increase with increasing molecular weight. Yet the melt

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A The boiling points of HCl, HBr, and HI increase with...

A The boiling points of $\mathrm{HCl}, \mathrm{HBr}$, and HI increase with increasing molecular weight. Yet the melting and boiling points of the sodium halides, $\mathrm{NaCl}$, $\mathrm{NaBr}$, and NaI, decrease with increasing formula weight. Explain why the trends are opposite. Describe the intermolecular forces present in each of these compounds and predict which has the lowest boiling point.

Key Concepts

Intermolecular Forces versus Ionic Bonds

Molecular compounds such as HCl, HBr, and HI are held together in the liquid phase largely by intermolecular forces, including dipole–dipole interactions and London dispersion forces. In contrast, the sodium halides are ionic compounds whose high melting and boiling points arise from strong electrostatic attractions between oppositely charged ions, known as ionic bonds. This fundamental difference in bonding leads to different trends in phase transition temperatures.

London Dispersion Forces and Polarizability

In the hydrogen halides, dispersion forces play a significant role, and these forces become stronger as the molecular weight and electron cloud size increase. The increased polarizability of the larger halogen atoms in HBr and HI results in stronger induced dipole-induced dipole attractions, thereby raising the boiling points with increasing molecular weight.

Lattice Energy and Ionic Sizes

For the sodium halides, the key factor is the lattice energy, which is the energy released when oppositely charged ions form a crystalline lattice. Although ionic bonds are inherently strong, the lattice energy is influenced by the sizes of the ions. As the halide ion increases in size (from Cl– to I–), the distance between the ions increases, reducing the lattice energy and consequently lowering the melting and boiling points despite the increase in formula weight.

Predicting the Lowest Boiling Point

Comparing the two sets of compounds, the hydrogen halides have boiling points determined by relatively weak intermolecular forces, whereas the sodium halides have much higher phase transition temperatures due to strong ionic bonds. Among the hydrogen halides, HCl has the lowest boiling point because it has the lowest molecular weight and the least polarizable halogen component, resulting in weaker dispersion forces compared to HBr and HI.

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