States of Matter

The particle model

The kinetic particle model treats matter as composed of particles — atoms, molecules or ions — in continuous motion. The state of a substance depends on two competing factors:

• Kinetic energy — proportional to temperature. Higher temperature = faster particle motion.
• Intermolecular forces — attractions between particles. Stronger forces favour closer packing (solid or liquid).

The types of force that hold particles together — ionic bonds, covalent bonds, van der Waals forces, hydrogen bonds — are covered in our chemical bonding guide. The stronger the forces, the higher the melting and boiling points.

For the atomic basis of these forces — why atoms attract each other — see our atomic structure guide (electron shells, electronegativity, ionisation energy).

Changes of state

A change of state occurs when energy input or removal shifts the balance between kinetic energy and intermolecular forces:

Critically, during a change of state, temperature does not change. Energy being supplied goes into breaking intermolecular forces (endothermic) rather than increasing kinetic energy. This energy is called latent heat:

• Specific latent heat of fusion (L_f) — energy per kg to melt a solid at its melting point. For water: 334 kJ/kg.
• Specific latent heat of vaporisation (L_v) — energy per kg to vaporise a liquid at its boiling point. For water: 2260 kJ/kg — the reason sweating is such an effective cooling mechanism.

Why substances have different melting and boiling points

The melting and boiling points of a substance reflect the strength and type of its interparticle forces:

• Ionic solids (e.g., NaCl) — strong electrostatic attractions between oppositely charged ions → high melting points (NaCl: 801 °C). They are brittle because displacing layers brings like charges adjacent, causing repulsion.
• Simple molecular solids (e.g., I₂, CO₂) — held together only by weak van der Waals forces → low melting points. Molecular mass determines the strength of dispersion forces.
• Hydrogen-bonded molecular liquids (e.g., H₂O, NH₃, HF) — hydrogen bonds are stronger than van der Waals forces but weaker than ionic/covalent bonds → intermediate boiling points, anomalously high for molecular mass. Water's boiling point (100 °C) is approximately 160 °C higher than would be expected from its molecular mass alone.
• Giant covalent structures (e.g., diamond, SiO₂) — covalent bonds throughout the lattice → extremely high melting points (diamond: ~3550 °C).
• Metallic structures — electrostatic attraction between delocalised electrons and positive ions → high melting points (iron: 1538 °C) and good electrical conductivity.

For how to calculate how many moles are involved in phase change problems, see our moles guide. For the reaction chemistry that often accompanies dissolution and phase changes, see our reaction rates guide.

Phase diagrams

A phase diagram plots the state of a pure substance as a function of temperature and pressure. Three key features:

• Triple point — the unique temperature-pressure combination at which solid, liquid and gas coexist in equilibrium. For water: 0.01 °C at 611.7 Pa.
• Critical point — above this temperature and pressure, the distinction between liquid and gas disappears and a supercritical fluid forms. For CO₂: 31.1 °C and 73.8 atm. Supercritical CO₂ is widely used as an industrial solvent in decaffeination and pharmaceutical extraction.
• Water's anomaly — the melting curve of water has a negative slope (melting point decreases slightly with pressure), the opposite of almost all other substances. This is because ice is less dense than water — ice floats.

Understanding phase behaviour at the molecular level requires knowing how intermolecular forces respond to temperature and pressure — topics covered in our chemical bonding guide and the covalent bond glossary entry.