8.3: Background - Density and Intermolecular Forces

Dispersion Force

Dispersion forces are caused by temporary dipoles within molecules. Electron density within a molecule shifts, momentarily causing one part of a molecule to have a partial negative charge (δ-). The region with less electron density then momentarily has a partial positive charge (δ+). The partially positive region then attracts the electron density of a neighboring molecule. As the electrons are constantly shifting, the dipoles change within the molecules. All molecules have electrons so all molecules can interact via dispersion forces.

The strength of dispersion forces depends on how easily the electron density shifts within a molecule. The technical name for this process is the polarizability of the molecule. The more electrons and the large the molecule, the more polarizable a molecule becomes. Thus, as the number of atoms or the size of the atoms of a molecule increases, the more dispersion forces are present in a molecule.

Dipole-Dipole Interaction

Polar molecules have permanent regions of partial positive (δ+) and partial negative (δ-) charge due to the electronegativity difference between the atoms. The dipoles of the molecules can then align so the partially positive region in one molecule interacts with the partially negative region in another molecule. This is shown in Figure \(\PageIndex{1}\) .

Due to the dipoles in polar molecules being permanent instead of temporary, this intermolecular force is stronger. This can be seen by comparing the boiling point of SiH₄ (silane) to HCl (hydrogen chloride). Both have a similar molar mass, 32.117 g/mol for silane and 36.46 g/mol for hydrogen chloride, so the dispersion forces are expected to be similar. However, the boiling point at 1 atm for silane is -185°C as compared to hydrogen chloride boiling point at 1 atm of -85°C. Thus, more energy is required to separate molecules of the polar hydrogen chloride as compared with the nonpolar silane molecules.

Hydrogen bonding

First and foremost, hydrogen bonding is not as strong as an ionic or covalent bond! The name is misleading. Hydrogen bonding (H-bonding) interactions are extra strong dipole-dipole interactions. They occur only in systems that contain an O-H, N-H, or H-F bond. The extreme dipole is causes by the large electronegativity difference between oxygen, nitrogen, and fluorine relative to hydrogen. The second factor that explains this extreme dipole is that the hydrogen atom has a very small valence shell, so when the electron density is shifted away from the hydrogen atom, the positively charged proton in the hydrogen atom’s nucleus is more exposed and is not shielded by inner electrons as happens in larger atoms.

The strength of the hydrogen bonding interaction can be seen in comparing the boiling points at 1 atm of hydrogen chloride and hydrogen fluoride. Hydrogen fluoride has a much lower molar mass than hydrogen chloride, so based on dispersion forces, the prediction would be that it would also have a lower boiling point. However, the boiling point of hydrogen fluoride is 20°C, which is much higher than -85°C for hydrogen chloride.

Ion-Dipole Interactions

Ion-dipole interactions occur when ions (which are charged) interact with polar molecules in solution. For example, when hydrogen chloride is added to water, it becomes hydrochloric acid, which has H₃O⁺ and Cl^- ions surrounded by water molecules. The water molecules are polar and have partially negative regions and partially positive regions. The partially negative regions of water interact with the positively charged H₃O⁺ (hydronium) ions. The partially positive regions of water interact with the negatively charged Cl^- (chloride) ions.