Classification of Solids

Matter around us commonly exists in three physical states — solid, liquid and gas — and this chapter builds a

quantitative picture of the two extremes, the solid state and the gaseous state, starting with how solids are

classified. Solids are characterised by a definite shape and volume, incompressibility, and a rigidity that comes

from strong, fixed forces holding the constituent particles at essentially unchanging positions. Chemists classify

solids into four broad types based on what the constituent particles are and what force holds them together at

their lattice points.

Molecular solids are built from discrete molecules held together only by comparatively weak intermolecular

forces — London dispersion forces, dipole-dipole forces, or hydrogen bonds. Because these binding forces are weak,

molecular solids are typically soft, have low melting points, and are poor conductors of electricity. Solid carbon

dioxide (dry ice, held by weak dispersion forces between non-polar $\text{CO}_2$ molecules), solid iodine, and ice

(held together by hydrogen bonds between polar $\text{H}_2\text{O}$ molecules) are classic examples.

Ionic solids are built from positive and negative ions held at fixed lattice positions by strong, non-directional

electrostatic (Coulombic) forces. This strong binding gives ionic solids high melting points and hardness, but they

are brittle (a small displacement brings like charges into contact and the lattice shatters) and conduct electricity

only when molten or dissolved, never in the solid state, because the ions cannot move freely through the rigid

lattice. Sodium chloride ($\text{NaCl}$) and magnesium oxide ($\text{MgO}$) are typical examples.

Covalent (network) solids are giant, three-dimensional structures in which every atom is joined to its

neighbours by strong, directional covalent bonds that extend throughout the entire crystal — the whole crystal is,

in effect, one giant molecule. This makes covalent solids extremely hard, with very high melting points, and (except

for graphite, whose delocalised electrons make it a special case) electrical insulators. Diamond, in which each

carbon atom is tetrahedrally bonded to four others, and quartz ($\text{SiO}_2$) are the standard examples.

Metallic solids consist of positively charged metal ion cores arranged in a regular lattice, immersed in a "sea"

of delocalised, mobile valence electrons that belong to the crystal as a whole rather than to any one atom. This

electron sea explains metals' characteristic properties: good electrical and thermal conductivity (the mobile

electrons carry charge and energy), malleability and ductility (the ion cores can slide past one another without

breaking the overall metallic bonding, since the electron sea simply flows around them), and a lustrous appearance.

Copper, iron and sodium are typical metallic solids.

Applying this classification: $\text{NaCl}$ is ionic (electrostatic attraction between $\text{Na}^+$ and

$\text{Cl}^-$); diamond is covalent/network (a continuous 3-D lattice of C-C covalent bonds); solid $\text{CO}_2$

(dry ice) is molecular (weak dispersion forces between discrete, non-polar $\text{CO}_2$ molecules); and copper is

metallic (ion cores in a delocalised electron sea). Recognising which of these four categories a given solid falls

into is the first step to predicting its physical behaviour — melting point, hardness, brittleness, and electrical

conductivity all follow directly from the nature of the binding force.