What do we mean by "states of matter"?

Everything around you — the air you breathe, the water you drink, the desk you're sitting at — is matter, and matter exists in three familiar forms: solid, liquid, and gas. (A fourth state, plasma, exists too — ionised gas found in things like lightning and stars — but at this level we focus on the three everyday states.) What decides which state a substance is in, at any given moment, is a tug-of-war between two things: how much kinetic energy the particles have (which makes them want to move and spread apart) and how strong the forces of attraction between the particles are (which pulls them together and holds them in place). Temperature controls the first; the identity of the substance controls the second.

The Kinetic Theory of Matter

The Kinetic Theory is the model chemists use to explain the behaviour of matter in all three states. It rests on a few core ideas:

1. All matter is made of tiny particles (atoms, molecules, or ions) that are in constant, random motion.
2. The particles possess kinetic energy, and the amount of kinetic energy they have increases as temperature increases — temperature is essentially a measure of the average kinetic energy of the particles in a sample.
3. There are forces of attraction between particles, and these forces try to pull particles together.
4. The state a substance is in depends on the balance between particle kinetic energy (which favours disorder and separation) and the attractive forces (which favour order and closeness).

Think of it like a tug-of-war: at low temperature, the attractive forces win, and particles are locked closely together — that's a solid. As temperature rises, the particles gain enough energy to partly overcome the attractive forces and slide past each other — that's a liquid. At even higher temperature, the particles have enough energy to escape the attractive forces almost entirely and fly off independently — that's a gas.

Properties of the three states, explained by particle behaviour

Solids

In a solid, particles are packed very close together in a fixed, regular arrangement (a lattice), held tightly by strong attractive forces. The particles don't move from place to place — they only vibrate about fixed positions. This is why solids have:

• A fixed shape and fixed volume (particles can't move past each other to change shape).
• High density (particles are packed tightly with little empty space).
• Incompressibility (there's almost no empty space left to squeeze out).

Liquids

In a liquid, particles are still close together (similar spacing to a solid) but the attractive forces are weaker relative to the particles' kinetic energy, so particles can slide and move past one another, though they can't easily escape each other's attraction entirely. This is why liquids have:

• No fixed shape — a liquid takes the shape of its container, because particles can flow.
• A fixed volume — particles are still held close enough together that they don't spread apart to fill all available space.
• Moderate density and low compressibility, similar to solids since spacing is similar.

Gases

In a gas, particles have enough kinetic energy to almost completely overcome the attractive forces between them. They move fast, independently, and in random straight lines until they collide with something, with huge distances between particles compared to their own size. This is why gases have:

• No fixed shape and no fixed volume — a gas expands to fill whatever container it's in.
• Low density (mostly empty space between particles).
• High compressibility (there's a lot of empty space to squeeze the particles into).

Changes of state and the energy involved

Changing from one state to another (a phase change) happens when enough energy is added or removed to shift the balance between kinetic energy and attractive forces. The names of these transitions:

• Melting: solid → liquid (energy absorbed, at the melting point)
• Freezing: liquid → solid (energy released, at the same temperature as the melting point)
• Boiling/vaporisation: liquid → gas (energy absorbed, at the boiling point)
• Condensation: gas → liquid (energy released)
• Sublimation: solid → gas directly, without passing through the liquid state (energy absorbed) — e.g. solid $CO_2$ ("dry ice") and iodine both do this at ordinary pressure.
• Deposition: gas → solid directly (energy released)

A crucial and commonly misunderstood point: during a phase change, temperature stays constant, even though energy (heat) is still being added or removed. Imagine heating a block of ice steadily: the temperature rises from, say, $-10°C$ up to $0°C$ — but then it stops rising and stays at exactly $0°C$ for a while, even as you keep heating it, until all the ice has melted. Only after all the ice is gone does the temperature (now of liquid water) start rising again. This is because, during melting, all the extra energy supplied is being used to break the attractive forces holding the solid lattice together (increasing potential energy between particles), rather than to speed the particles up (which would increase kinetic energy and therefore temperature).

The energy absorbed or released during a phase change (with no temperature change) is called latent heat: specifically latent heat of fusion (melting/freezing) and latent heat of vaporisation (boiling/condensing). "Latent" literally means "hidden" — the energy is going into the substance, but you can't detect it as a temperature rise on a thermometer, because it's being spent overcoming intermolecular forces rather than speeding particles up.

Why boiling and melting points differ between substances

The strength of the attractive forces between particles determines how much energy is needed to change state, and therefore determines the melting and boiling points. Substances with strong forces between their particles (e.g. ionic compounds like $NaCl$, held together by strong electrostatic attraction, or metals) have high melting and boiling points, because a lot of energy is needed to overcome those strong forces. Substances with weak forces between their particles (e.g. small covalent molecules like $O_2$ or $CO_2$, where only weak intermolecular forces exist between separate molecules) have low melting and boiling points.

Worked example: table salt, $NaCl$, melts at $801°C$, while oxygen, $O_2$, melts at $-218°C$. Both are made of small particles, so why the huge difference? In solid , the particles are and ions held together by strong electrostatic (ionic) attraction throughout the whole lattice — a lot of energy is needed to break this apart. In solid , the particles are molecules, and although the atoms each molecule are strongly covalently bonded to each other, the forces separate molecules (weak van der Waals forces) are what must be overcome to melt the solid — and these are far weaker, so only a little energy (a low temperature) is needed.

Evaporation vs boiling: a common point of confusion

Both are examples of liquid turning to gas, but they're not the same thing:

• Boiling happens throughout the whole liquid, only at the boiling point, and only when the liquid's vapour pressure equals the surrounding atmospheric pressure.
• Evaporation happens only at the liquid's surface, can occur at any temperature below the boiling point, and happens because a small fraction of particles at the surface always have enough kinetic energy (due to the natural spread of particle energies) to escape into the gas phase, even though the bulk of the liquid is nowhere near boiling.

This is why a puddle of water evaporates and disappears on a warm day even though it's nowhere near $100°C$ — a few fast-moving surface molecules keep escaping little by little.

Why This Matters for Your Exams

Exam questions on this topic frequently ask you to explain — in terms of particle arrangement, movement, and forces of attraction — why a substance behaves a certain way, rather than just to state a property. Always answer in terms of the Kinetic Theory: mention particle spacing, particle movement, and the relative strength of attractive forces versus kinetic energy. The single most commonly tested "trick" idea is that temperature stays constant during a phase change — if you're asked why a heating curve has flat sections, the answer is always that the energy is being used to overcome intermolecular/interparticle forces (changing potential energy), not to speed particles up (which would change kinetic energy and thus temperature). Also be ready to distinguish evaporation from boiling clearly, since exams like to test this contrast directly.