A molecular solid is a solid made up of individual molecules held together by weak attractive forces between them, rather than by the strong bonds that hold atoms together within each molecule. Ice, dry ice, sugar, and solid iodine are all molecular solids. Because the forces holding the molecules in place are relatively weak, molecular solids tend to be soft, have low melting points, and don’t conduct electricity.
Understanding molecular solids comes down to one key distinction: the bonds inside each molecule are strong, but the forces between neighboring molecules are much weaker. When you melt a molecular solid, you’re not breaking molecules apart. You’re just giving them enough energy to slip past each other.
How Molecules Hold Together in a Solid
At low temperatures, the molecules in a molecular solid lock into a rigid, repeating structure. What keeps them there are intermolecular forces, sometimes called van der Waals forces. These come in three main types, ranging from weakest to strongest.
- London dispersion forces are the weakest and most universal. Every molecule experiences them. They arise because electrons are constantly moving around within a molecule, creating tiny, fleeting imbalances in charge. These momentary imbalances cause neighboring molecules to develop matching imbalances, creating a brief mutual attraction. Even completely nonpolar molecules like oxygen or methane experience dispersion forces. Larger molecules with more electrons generate stronger ones, which is why bigger molecules tend to have higher melting points.
- Dipole-dipole interactions occur in molecules where charge is unevenly distributed. One end of the molecule carries a slight positive charge and the other a slight negative charge. Neighboring molecules orient themselves so that opposite charges face each other, creating a steady attraction. The strength of this interaction depends on how polar the molecule is and how close the molecules are to each other.
- Hydrogen bonds are the strongest of the three. They form when a hydrogen atom bonded to oxygen, nitrogen, or fluorine interacts with a lone pair of electrons on a nearby molecule’s oxygen, nitrogen, or fluorine. Despite the name, hydrogen bonds are not true chemical bonds. They’re simply very strong dipole-dipole interactions. Water, for example, owes many of its unusual properties to hydrogen bonding.
Most molecular solids are held together by some combination of these forces. A nonpolar molecule like solid methane relies entirely on London dispersion forces. A polar molecule like solid hydrogen chloride adds dipole-dipole interactions. Ice benefits from all three, with hydrogen bonds doing most of the heavy lifting.
Physical Properties of Molecular Solids
The weakness of intermolecular forces gives molecular solids a distinctive set of physical properties. They tend to be soft and often crumble or deform easily. Their melting points are low compared to other types of solids, though the range is wide depending on molecular size and the types of intermolecular forces involved.
To put some numbers on it: solid chlorine melts at -101.5 °C, dry ice (solid carbon dioxide) sublimes directly into gas at -78.5 °C, and solid iodine melts at 113.7 °C. Even at the high end, these temperatures are modest compared to ionic solids like table salt (801 °C) or covalent network solids like diamond, which requires thousands of degrees to break down. The pattern within molecular solids is consistent: heavier molecules and stronger intermolecular forces mean higher melting points.
Molecular solids are poor conductors of electricity. Their electrons are tightly held within individual molecules, with no free-moving charged particles to carry a current. This stands in contrast to metals, where electrons flow freely, or ionic solids, which conduct when melted or dissolved.
Ice as a Molecular Solid
Ice is one of the most familiar and instructive examples. In ordinary ice (called ice Ih), each water molecule sits in a hexagonal crystal lattice. Every oxygen atom is surrounded by four neighboring oxygen atoms arranged in a tetrahedron, with oxygen-to-oxygen distances of about 2.75 angstroms. Two hydrogen bonds point outward from each molecule toward neighbors, and two point inward from neighbors toward it.
This tetrahedral arrangement creates an open, spacious structure with a lot of empty space between molecules. That’s why ice is less dense than liquid water and floats. When ice melts, some of those hydrogen bonds break, the rigid structure collapses, and molecules pack more tightly together. It’s a direct consequence of the geometry that hydrogen bonding imposes on the solid.
How They Differ From Other Solids
The easiest way to understand molecular solids is to contrast them with the other major types of solids. In a covalent network solid like diamond or quartz, atoms are connected by strong covalent bonds that extend throughout the entire structure. There are no distinct molecules. To melt diamond, you’d need to break those covalent bonds directly, which requires extreme temperatures. Diamond is also the hardest known natural material, precisely because of that continuous bonding network. Molecular solids, by comparison, are soft and brittle because you only need to overcome weak intermolecular attractions to deform or melt them.
Ionic solids like table salt are held together by the electrostatic attraction between positive and negative ions. These forces are much stronger than intermolecular forces, giving ionic solids high melting points and making them hard and brittle. Unlike molecular solids, ionic solids can conduct electricity when dissolved in water because their ions become free to move.
Metallic solids share electrons across a “sea” of delocalized electrons, which makes them excellent conductors and gives them the ability to bend without breaking. Molecular solids share none of these properties.
Solubility Patterns
Whether a molecular solid dissolves in a given liquid follows a simple guideline: like dissolves like. Polar molecular solids tend to dissolve in polar solvents like water. Sugar dissolves in water because both are polar. Ethanol, a small polar molecule, is also freely soluble in water.
Nonpolar molecular solids dissolve in nonpolar solvents. Hydrocarbons like wax dissolve in hexane or other organic solvents but won’t dissolve in water to any meaningful degree. Small polar molecules sometimes dissolve in both polar and nonpolar solvents, which is why substances like acetone and ethanol can mix with water and with organic solvents.
Common Examples
Molecular solids are everywhere in daily life. Ice, sugar, naphthalene (the compound in mothballs), solid carbon dioxide (dry ice), and chocolate are all molecular solids. Most organic compounds form molecular solids when cooled enough, including many medications in pill form. The crystals of iodine you might see in a chemistry lab, with their distinctive purple-gray sheen, are a classic example of a nonpolar molecular solid held together almost entirely by London dispersion forces.
What unites all of these is the same basic architecture: intact, individual molecules arranged in a crystal lattice, held in place by intermolecular forces that are strong enough to maintain structure at low temperatures but weak enough to break apart with relatively modest amounts of heat.