Molecules, Bonds & States of Matter

Steel holds a bridge up, water flows, air fills a room. All three are made of atoms; what differs is how strongly and how closely those atoms hold on to each other.

Layers of structure#

Nuclear forces between protons and neutrons (together, nucleons) shape the nucleus. Electromagnetic forces between electrons and nuclei shape the atom. Forces between atoms build molecules, and forces between molecules give a material the properties we see and touch.

The force between two atoms#

Plot the potential energy of a pair of atoms against their separation r, taking U = 0 when they are far apart. As they approach, U falls below zero: the force is attractive. At a particular separation r₀, U is lowest: the force is zero and the atoms can rest in equilibrium. Closer still, U rises steeply: the force becomes strongly repulsive.

A molecule forms when its atoms settle into the arrangement of lowest total potential energy. Forces between molecules follow the same pattern: weak attraction far away, growing as they approach, zero at r₀, repulsion closer in.

Bonds#

Atoms bind mainly through electrostatic attraction between electrons and nuclei. Two kinds of bond are especially common:

Solids, liquids and gases#

Two molecules at separation r₀ are in equilibrium. Pull them slightly apart and attraction pulls them back; push them together and repulsion pushes them apart. So a displaced molecule oscillates about r₀.

• Solid: molecules sit very nearly at their equilibrium separations and vibrate with tiny amplitude about fixed positions, so the solid keeps its shape.
• Liquid: the average separation is a little larger, the attraction weaker, and molecules move freely through the whole volume.
• Gas: molecules are far apart and the forces between them are very weak.

Kinds of solid#

In many solids the equilibrium positions form a regular three-dimensional pattern, a crystal; the positions are lattice points. Crystalline solids are classified by what holds them together:

Molecular solids

Covalent molecules held together by weaker forces between molecules: dipole–dipole bonds between polar molecules (like water, whose positive and negative charge centres don’t coincide) and van der Waals bonds between nonpolar ones (H₂, O₂, Cl₂). Usually soft, with low melting points; poor conductors.

Ionic solids

Positive and negative ions on the lattice points, bound by strong electrostatic forces. Hard, with fairly high melting points; poor conductors.

Covalent solids

Atoms joined to their neighbours by shared electrons in a network extending through the whole crystal. In diamond each carbon atom bonds to four others. Very hard, high melting point, poor conductors.

Metallic solids

Positive ions on the lattice points, with the detached electrons roaming the whole solid like a gas. Those free electrons make metals excellent conductors.

Amorphous (glassy) solids#

Some solids have only local order: small groups of 4–5 molecules bond together, but these groups are arranged randomly, like the molecules of a liquid. The forces are much stronger than in a liquid, so the material doesn’t flow. Glass (strong Si–O–Si bonds that don’t extend far) is the classic example.