A substance’s melting point is determined primarily by how strongly its particles attract each other and how efficiently they pack together in a solid. Stronger attractive forces and tighter packing both mean more energy is needed to break the solid apart, which translates to a higher melting point. But several other factors, from molecular shape to pressure to impurities, also play a role.
Intermolecular Forces: The Biggest Factor
The attractive forces between molecules are the single most important factor in setting a substance’s melting point. These forces exist on a clear hierarchy, and each step up the ladder corresponds to dramatically higher melting temperatures.
- Ionic bonds are the strongest. Substances held together by the electrostatic pull between positive and negative ions require enormous energy to melt. Table salt, for example, melts at 801 °C.
- Hydrogen bonds form when hydrogen is attached to oxygen, nitrogen, or fluorine. These bonds make molecules “sticky,” raising the melting point well above what you’d expect based on molecular weight alone. This is why water (with extensive hydrogen bonding) is a liquid at room temperature while similar-sized molecules without hydrogen bonds are gases.
- Dipole-dipole forces occur between molecules that have a permanent uneven charge distribution. These attractions are weaker than hydrogen bonds but still pull melting points upward compared to nonpolar molecules of similar size.
- Dispersion forces (also called London forces) are the weakest. They exist in all molecules but are the only attractive force in nonpolar substances like hydrocarbons. Dispersion forces grow stronger with molecular size, which is why small hydrocarbons like methane are gases while larger ones like paraffin wax are solids.
Molecular Size and Weight
Within any family of similar compounds, larger molecules have higher melting points. This pattern shows up clearly in straight-chain hydrocarbons (alkanes): as you add carbon atoms to the chain, both the melting and boiling points rise in a regular, predictable way. Methane (one carbon) melts at -182 °C, while tetradecane (fourteen carbons) melts around 6 °C.
The reason is straightforward. Bigger molecules have more surface area, which gives dispersion forces more area to act across. Think of it like Velcro: a bigger strip holds more firmly than a smaller one, even though each individual hook is equally weak.
Molecular Shape and Symmetry
Two molecules with the same chemical formula can have very different melting points depending on their shape. Compact, symmetrical molecules pack more efficiently into a crystal, like uniform bricks stacking neatly into a wall. That tight packing creates stronger overall attractions and a higher melting point.
This shows up clearly in geometric isomers. Trans isomers, where groups sit on opposite sides of a double bond, tend to be more symmetrical and pack tightly, giving them higher melting points. Cis isomers, with groups on the same side, have a bent shape that prevents close packing, resulting in lower melting points. The classic example is fumaric acid (trans, melting point 287 °C) versus maleic acid (cis, melting point 135 °C). Same atoms, same bonds, but a 150-degree difference in melting point purely because of shape.
Branching in carbon chains has a similar effect, though the outcome depends on whether the branching makes the molecule more or less symmetrical. A highly branched molecule that becomes more spherical can actually pack better and melt higher than its straight-chain counterpart.
Crystal Structure and Lattice Packing
When a substance freezes, its particles arrange into a repeating three-dimensional pattern called a crystal lattice. How stable that lattice is directly affects the melting point. For ionic compounds, the key factors are the distances between ions in the crystal and the energy released when the solid forms from its elements (the formation enthalpy). Shorter distances between ions mean stronger electrostatic pull and a higher melting point, which is why lithium fluoride (small ions, short distance) melts at a higher temperature than cesium iodide (large ions, long distance).
For molecular solids, the efficiency of packing matters enormously. Molecules that fit together with minimal wasted space form more stable crystals. This is why symmetry plays such an outsized role: a perfectly spherical or cylindrical molecule leaves fewer gaps in the lattice than an oddly shaped one.
Impurities Lower the Melting Point
Adding even a small amount of a foreign substance to a pure solid will lower its melting point, a phenomenon called melting point depression. The mechanism comes down to disorder. An impure solid is a patchwork at the microscopic level, with tiny pure regions of each component scattered through the bulk material. When heated, the component with the lower melting point begins to melt first in its own pure pockets, and the resulting liquid is more disordered (higher entropy) than a pure liquid would be. That extra disorder makes melting thermodynamically favorable at a lower temperature.
This principle is useful in chemistry labs for checking whether a substance is pure. A pure solid melts sharply over a narrow range of 1 to 2 °C. An impure solid melts at a lower temperature and over a wider range, sometimes spanning 5 °C or more. The wider and lower the range, the more impurity is present. Road salt works on the same principle: mixing salt into ice lowers its melting point, which is why icy roads clear up after salting even when the air temperature is below freezing.
Pressure Effects
Pressure usually raises the melting point of a substance. Squeezing a solid makes it harder for molecules to move apart into the less-ordered liquid state, so more heat is needed. For most materials, the solid is denser than the liquid, and increasing pressure favors the denser phase, keeping the substance solid to higher temperatures. Carbon dioxide follows this typical pattern.
Water is a famous exception. Ice is less dense than liquid water (which is why ice floats), so increasing pressure actually favors the denser liquid phase and lowers the melting point. At 987 atmospheres of pressure, ice melts at -9 °C instead of the usual 0 °C. Antimony and bismuth share this unusual behavior, but for the vast majority of substances, more pressure means a higher melting point.
Under everyday conditions, pressure changes are too small to noticeably shift melting points. This effect only becomes significant at very high pressures, measured in hundreds or thousands of atmospheres.
Ionic vs. Covalent vs. Metallic Solids
The type of bonding within a solid sets the baseline range for its melting point. Ionic solids, held together by the attraction between charged ions, typically melt between a few hundred and over a thousand degrees Celsius. Metals, where positive ion cores sit in a “sea” of shared electrons, vary widely but many melt at high temperatures (iron at 1,538 °C, tungsten at 3,422 °C). Molecular solids, held together only by the weaker intermolecular forces described above, tend to have the lowest melting points. Most organic compounds melt below 300 °C, and many below room temperature.
Network covalent solids sit at the extreme high end. In these materials, every atom is connected to its neighbors by strong covalent bonds extending throughout the entire crystal. Diamond (carbon atoms bonded in a three-dimensional network) doesn’t truly melt under normal pressure but converts to graphite around 3,500 °C. Quartz (silicon dioxide) melts at about 1,713 °C. These materials resist melting because breaking them apart means breaking actual covalent bonds, not just overcoming weak intermolecular attractions.