Molecular solids are crystalline materials composed of discrete, neutral molecules. These molecules are held together in a solid structure by relatively weak electrostatic forces, known collectively as intermolecular forces (IMFs). This is in contrast to substances like table salt or metals, which are bound by much stronger ionic or metallic bonds. The forces holding the molecules together in the solid state are significantly weaker, which dictates many of the solid’s observable properties.
The Universal Force: London Dispersion
The most fundamental attractive force existing between all molecules, whether polar or nonpolar, is the London Dispersion Force (LDF). This force arises from the constant, random movement of electrons within a molecule, which can momentarily result in an uneven distribution of charge, creating a fleeting, instantaneous dipole moment. This temporary dipole then influences the electron cloud of a neighboring molecule, inducing a corresponding dipole. The resulting attraction is the LDF. Though individually weak, the cumulative effect allows nonpolar substances, like solid oxygen or dry ice, to solidify at low temperatures.
The strength of the London Dispersion Force is heavily dependent on two factors: the size and the shape of the molecule. Larger molecules generally possess stronger LDFs because they have a greater number of electrons, and the electron cloud is more easily distorted, a property called polarizability. Furthermore, molecules with a greater surface area can achieve more extensive contact with neighbors, maximizing the total attraction and leading to higher solidification temperatures.
Polar Attraction: Dipole-Dipole Interactions
The dipole-dipole interaction occurs exclusively between molecules that possess a permanent separation of charge, known as a net dipole moment. These molecules are classified as polar because of an uneven sharing of electrons between atoms due to differences in electronegativity. Examples include hydrogen chloride (HCl) and sulfur dioxide (\(\text{SO}_2\)).
The positive pole of one polar molecule is electrostatically attracted to the negative pole of an adjacent polar molecule. This permanent, directional attraction is stronger than LDF for molecules of comparable size. The molecules in the solid state will align themselves to maximize this attraction, forming an ordered lattice structure.
Dipole-dipole forces are typically moderate in strength, ranging from 5 to 20 kilojoules per mole (\(\text{kJ/mol}\)). While stronger than LDF for molecules of similar mass, this attraction is still considerably weaker than the actual covalent bonds that hold the atoms together within the molecule itself.
The Strongest Intermolecular Force: Hydrogen Bonding
Hydrogen bonding represents a particularly strong and specific type of dipole-dipole interaction. It is an unusually powerful intermolecular attraction that occurs only under precise conditions. The force requires a hydrogen atom to be directly bonded to one of the three most electronegative non-metals: Nitrogen (N), Oxygen (O), or Fluorine (F).
This highly polarized covalent bond draws shared electrons toward the N, O, or F atom, leaving the hydrogen atom with a significant partial positive charge. This partially positive hydrogen is then strongly attracted to a lone pair of electrons on a neighboring N, O, or F atom. Water (\(\text{H}_2\text{O}\)) is the most familiar example, where each molecule can form up to four hydrogen bonds with its neighbors.
In ice, this strong directional force compels water molecules to arrange into a highly organized, open, three-dimensional hexagonal lattice structure. This arrangement creates empty space within the crystal. The stability and geometry of these hydrogen bonds are responsible for the unique property that solid ice is less dense than liquid water, allowing it to float.
Physical Consequences of Weak Intermolecular Forces
The relatively weak nature of all intermolecular forces—LDF, dipole-dipole, and hydrogen bonds—has profound effects on the bulk properties of molecular solids. Since little energy is needed to overcome the attractions, molecular solids are characterized by low melting and boiling points. For instance, dry ice (solid \(\text{CO}_2\)) sublimes at \(-78.5^\circ \text{C}\).
These solids also tend to be soft and easily deformed because the molecules can be pushed out of position without breaking the strong covalent bonds inside the molecule. The weak IMFs allow the layers of molecules to slide past one another relatively easily. This softness and low melting point are direct results of the forces holding the solid together being mere fractions of the strength of a true chemical bond.
Furthermore, molecular solids are poor electrical conductors, meaning they act as insulators. This is because the electrons are tightly held within the covalent bonds of each discrete molecule and are not free to move throughout the solid. Unlike metallic solids, molecular solids lack any free-moving charged particles necessary to carry an electrical current.