An equilibrium phase transition occurs when the chemical potentials of the two phases are equal. The chemical potential is also the molar Gibbs free energy $g\equiv h-Ts$, with molar enthalpy $h$, temperature $T$ and molar entropy $s$, so we have
$$h_\mathrm{liquid}-T_\mathrm{melting}s_\mathrm{liquid}=h_\mathrm{solid}-T_\mathrm{melting}s_\mathrm{solid},$$
or
$$T_\mathrm{melting}=\frac{\Delta h}{\Delta s},$$
referring to the difference in enthalpies and entropies between the two phases. Thus, better bonding in the solid state promotes a higher melting temperature, and greater entropy (e.g., more accessible molecular configurations) in the liquid state promotes a lower melting temperature. This tradeoff arises everywhere: Nature prefers* both strong bonding and high entropy, and the temperature mediates whether the former or the latter predominates. As low temperatures, we encounter solids (strong bonding), and at high temperatures, we encounter gases (high entropy).
*The reason is that high entropy offers many possible arrangements, so we're more likely to see one of them. Strong bonding releases energy that heats the rest of the universe, raising its entropy correspondingly; the same reasoning applies.
