The bonds we have just discussed (covalent, ionic and metallic) are all examples of intramolecular forces: the forces inside a molecule that keep it together. However, molecules commonly interact with each other. For instance, the formation of solids is due to the compounds being within close proximity to each other. What is holding these molecules together?

Dipole-Dipole Interactions and Hydrogen Bonding

Like the ends of a magnet, opposite charges attract, while similar charges repel. As a result, attraction forces between molecules are naturally present in those with charge. Recall that hydrogen chloride, a polar covalent molecule, experiences partial charges on its constituents. In this case, chlorine is more electronegative and is partially negative, given the closeness of the electron to the atom. On the other hand, hydrogen is partially positive because the electron is further away (see Figure 3.5.1).

When a molecule of hydrogen chloride is surrounded by another molecule of hydrogen chloride, it will begin to align according to attraction forces. Opposite ends will neighbour each other, as seen in Figure 3.5.2.

This is known as a dipole-dipole interaction – where molecules experience attraction and repulsion forces due to a permanent charge. Hydrogen bonds are a special kind of dipole-dipole interaction, as hydrogen exhibits unusually strong intermolecular forces when paired with a highly electronegative atom such as chlorine or oxygen (see Figure 3.5.3).

Intermolecular forces allow us to understand why molecules have a tendency to stay together (see Figure 3.5.3). It also explains the melting and boiling points of different compounds. Molecules that are closely attracted to one another require higher kinetic energy and, therefore, higher temperature to separate and undergo a phase change. The greater the strength of the attraction and the more interactions that are allowed to occur, the greater the energy requirement. This is why ionic compounds (particularly salts) generally have high melting points. For reference, sodium chloride has a melting point of 802°C and a boiling point of 1465°C!

London Dispersion Forces

Under this theory, we should expect that for non-polar covalent molecules that don’t hold permanent dipoles, melting and boiling points should be the same due to a lack of partial charge. However, as Table 3.5.1 demonstrates, boiling points do change with respect to the length of the molecule.

All molecules, regardless of dipoles, are subject to dispersion forces (also known as London forces). These occur due to electrons continuously shifting within molecules. These electrons may occasionally generate partial charges, which induce electrons in other atoms to do the same, generating a slightly attractive force between molecules. As the surface area of a molecule increases, so does the occurrence of dispersion forces. As a result, longer-chain alkanes (listed in Table 3.5.1) possess higher boiling points.

Those are the three kinds of intermolecular forces: hydrogen bonding, dipole-dipole interactions and London dispersion forces. You can see them listed in Table 3.5.2 and their relative strengths. In addition to dictating boiling and melting points within molecules, they also play an important part in the solubility of molecules within water.