Melting point is a distinct physical property of a substance, defined as the temperature at which its solid and liquid phases exist in equilibrium. The transition from a solid, which has an organized crystal structure, to a liquid requires an input of thermal energy. This energy is needed to overcome the attractive forces that hold the constituent particles—whether they are atoms, ions, or molecules. The melting temperature is therefore a direct measure of the strength of these attractive forces.
The Impact of Intermolecular Forces
For molecular substances, the melting point is determined by the relatively weak forces between molecules, known as intermolecular forces (IMFs). These forces are significantly weaker than the chemical bonds holding the atoms together within each molecule, which is why molecular solids tend to have lower melting points. Breaking these attractions allows the molecules to move past one another, transitioning into the liquid state.
The weakest of these attractions are the London Dispersion Forces (LDF), which arise from momentary, random fluctuations in electron distribution that create temporary dipoles in nonpolar molecules. A slightly stronger type is the dipole-dipole interaction, which occurs between polar molecules that possess a permanent separation of positive and negative charge. Both LDF and dipole-dipole forces are collectively referred to as van der Waals forces.
The strongest of the IMFs is hydrogen bonding, a specialized, strong dipole-dipole interaction. This occurs when a hydrogen atom is covalently bonded to a highly electronegative atom like nitrogen, oxygen, or fluorine. The resulting attraction requires significant energy to break, leading to substances like water having a much higher melting point than similarly sized molecules that lack this bonding capability.
The Influence of Chemical Bond Strength
In contrast to molecular solids, substances like salts and network solids possess extremely high melting points because the attractive forces holding them together are strong chemical bonds. Melting these materials requires breaking the actual ionic or covalent bonds that form the entire structure, which demands substantially greater thermal energy. This is why table salt (sodium chloride) melts at 801 degrees Celsius, while ice melts at zero degrees Celsius.
Ionic compounds, such as sodium chloride, are held together by strong electrostatic forces between positively and negatively charged ions arranged in a crystal lattice. The energy required to break this lattice, called the lattice energy, is a major factor in determining the melting point. This energy is directly proportional to the charge of the ions and inversely proportional to the distance between their nuclei, meaning smaller, more highly charged ions result in a higher melting point.
Covalent network solids, such as diamond or quartz, represent the extreme end of high melting points. In these structures, every atom is linked to its neighbors by a continuous network of strong covalent bonds. Melting such a substance requires breaking numerous chemical bonds, which is why diamond’s melting point is estimated to be near 4,000 degrees Celsius.
How Molecular Size and Structure Affect Melting
Beyond the type of force, the size and structure of a substance’s constituent particles influence its melting point. For molecular solids, a larger molecule generally has a higher melting point due to the increased strength of its London Dispersion Forces. As the size of a molecule increases, it possesses more electrons, which enhances the potential for temporary dipoles and strengthens the attractive forces between neighboring molecules.
This effect is demonstrated when comparing hydrocarbons, where a long-chain molecule like candle wax has a higher melting point than a short-chain molecule like methane gas. However, the molecule’s shape is just as important as its mass. Highly symmetrical molecules are capable of packing together more efficiently into a stable, ordered crystal lattice.
This efficient packing maximizes the number of attractive interactions between particles, whether they are IMFs or ionic forces. Consequently, a highly symmetrical molecule, even if it is a structural isomer of a less symmetrical one, will require more energy to disrupt its compact solid arrangement. The increased stability of this efficient packing results in a higher melting temperature for the symmetrical compound.