Why Do Melting and Boiling Points Increase in a Homologous Series?

Melting and boiling points generally increase along a homologous series because molecular size and the number of electrons increase. This strengthens London dispersion forces, so more energy is required to separate the molecules.

This is part of IB Chemistry Structure 3.2 and applies at both SL and HL.

Members of a homologous series have the same functional group and similar chemical properties, but successive members usually differ by a $\text{CH}_2$ unit. For example, methane, ethane, and propane contain progressively longer carbon chains.

Adding each $\text{CH}_2$ unit increases the number of electrons and enlarges the electron cloud. The electron cloud becomes more polarizable, meaning temporary dipoles form more readily. These temporary dipoles induce dipoles in neighbouring molecules, producing stronger London dispersion forces.

Longer molecules can also have a larger contact area with neighbouring molecules. Consequently, more intermolecular attractions act between them. More thermal energy is therefore needed to overcome these forces during melting or boiling.

The boiling-point trend is usually smooth, but melting points can show irregularities. Molecules that pack more efficiently into a crystal lattice may have unexpectedly high melting points; chain shape, symmetry, and odd-even carbon-number effects can influence packing.

A common misconception is that stronger covalent bonds cause the increase. During melting and boiling, covalent bonds within molecules are not broken; only intermolecular forces are overcome.

For an IB explain question, link the full causal chain: increasing chain length $\rightarrow$ more electrons and greater polarizability $\rightarrow$ stronger London dispersion forces $\rightarrow$ more energy required. For melting points, add that crystal packing may produce deviations from the general trend.