In atomic solids, individual atoms are held in place by London forces. The noble gases are the only atomic solids known to form.
In molecular solids, lattices composed of molecules are held in place by London forces, dipole–dipole forces, and hydrogen bonding. Solid methane and water are examples of molecular solids.
In ionic solids, lattices composed of ions are held together by the attraction of the opposite charges of the ions. These crystalline solids tend to be strong, with high melting points because of the strength of the intermolecular forces. NaCl and other salts are examples of ionic solids. Figure 12.3 shows the lattice structure of NaCl. Each sodium cation is surrounded by six chloride anions, and each chloride anion is surrounded by six sodium cations.
In metallic solids, metal atoms occupying the crystal lattice are held together by metallic bonding. In metallic bonding, the electrons of the atoms are delocalized and are free to move throughout the entire solid. This explains electrical and thermal conductivity, as well as many other properties of metals.
In covalent network solids, covalent bonds join atoms together in the crystal lattice, which is quite large. Graphite, diamond, and silicon dioxide (SiO2) are examples of network solids. The crystal is one giant molecule.