States of Matter

Particle theory says all matter is made of tiny particles in constant motion. The three states differ in HOW the particles are arranged, HOW they move, and HOW STRONGLY they attract each other.

Solid

• Arrangement. Regular, tightly packed pattern (a lattice).
• Movement. Particles vibrate about fixed positions — they cannot move from place to place.
• Forces. Strong forces of attraction between every neighbour.
• Bulk properties. Fixed shape, fixed volume, cannot flow, cannot be compressed (no empty space to compress).

Liquid

• Arrangement. Close together but irregular / random — no long-range order.
• Movement. Particles slide past one another; some can escape from the surface at any temperature (evaporation).
• Forces. Moderate — strong enough to keep particles in contact, weak enough to let them move.
• Bulk properties. Fixed volume, takes the shape of the container, can flow, very hard to compress.

Gas

• Arrangement. Far apart, with large empty spaces between particles.
• Movement. Rapid, random straight-line motion in all directions until particles collide.
• Forces. Very weak / essentially negligible at IGCSE level.
• Bulk properties. No fixed shape, no fixed volume, spreads out to fill any container, easily compressed (because of all the empty space).

Mark-scheme tip. Edexcel awards marks for THREE keywords in a particle-arrangement answer: arrangement (close / far apart), movement (vibrate / slide / random), and forces (strong / moderate / weak). Hit all three for every state — examiner reports note that many candidates describe arrangement and movement but forget forces.

Six interconversions show up in 4CH1:

From → To	Process name	Energy flow
Solid → Liquid	Melting	Energy IN (forces broken)
Liquid → Solid	Freezing	Energy OUT (forces form)
Liquid → Gas (surface, any T)	Evaporation	Energy IN
Liquid → Gas (bulk, at bp)	Boiling	Energy IN
Gas → Liquid	Condensation	Energy OUT
Solid → Gas / Gas → Solid	Sublimation / deposition	Energy IN / OUT

Energy story.

• When forces are broken (melting, evaporation, boiling, sublimation), energy must be SUPPLIED. The particles gain potential energy as they pull apart against the attractions.
• When forces form (freezing, condensation, deposition), energy is RELEASED to the surroundings.

Examples of sublimation. Solid carbon dioxide (dry ice) sublimes directly to CO₂ gas at -78 °C without melting. Iodine crystals sublime when warmed gently in a test-tube. Ammonium chloride sublimes and re-deposits as a white solid on a cooler part of the apparatus — this is what produces the white ring in the NH₃ + HCl diffusion experiment.

Common 4CH1 question style. A graph or paragraph describes water heating from -10 °C to 110 °C. You must (a) label each region of the curve with the state(s) present, (b) name the change of state at each plateau, (c) explain why the temperature is constant during the plateau (energy used to overcome forces of attraction, not to raise kinetic energy).

A heating curve plots temperature against time as a substance is heated steadily from solid through to gas. It has FIVE regions for water:

1. Solid warming — sloping line going up. Particles gain kinetic energy, vibrate harder.
2. Melting plateau at 0 °C — horizontal line. ALL energy supplied goes to overcoming the forces of attraction between water particles. Kinetic energy (and so temperature) stays constant until all the solid has melted.
3. Liquid warming — sloping line going up. Particles gain kinetic energy, move faster.
4. Boiling plateau at 100 °C — horizontal line. Energy is used to break the remaining forces between particles so they can fully separate as gas.
5. Gas warming — sloping line going up.

A cooling curve is the reverse: gas → liquid → solid, with plateaus at the condensation and freezing points where energy is RELEASED to the surroundings (and the thermometer reads constant).

Why is the temperature constant on a plateau?

Temperature is a measure of the mean kinetic energy of particles. During a state change, the supplied energy increases the potential energy of the particles (pulling them apart against the forces of attraction) but does NOT increase their kinetic energy. So the temperature stays constant until the change is complete.

Edexcel pitfall. Candidates frequently write "the particles stop moving on the plateau" or "the heat is wasted". Neither is true. Particles still move; the heat is used to overcome forces of attraction. Use those exact words.

Diffusion = the net movement of particles from a region of higher concentration to a region of lower concentration as a result of their random motion. The result is a uniform concentration throughout.

Two factors set the rate:

1. Temperature. Higher T → particles have more kinetic energy → move faster → diffuse faster.
2. Relative molecular mass (Mr). At the same temperature, all particles have the same average kinetic energy ($\tfrac{1}{2} m v^2$ is fixed). So smaller $m$ → larger $v$ → lighter particles diffuse faster.

Classic NH₃ + HCl tube experiment.

A long horizontal glass tube has cotton wool soaked in concentrated ammonia (NH₃, Mr = 17) at one end and cotton wool soaked in concentrated hydrochloric acid (HCl, Mr = 36.5) at the other end. After a few minutes a white ring of solid ammonium chloride appears inside the tube:

$\text{NH}_3(g) + \text{HCl}(g) \rightarrow \text{NH}_4\text{Cl}(s)$

The ring forms closer to the HCl end because NH₃ is lighter and diffuses faster, so it travels further in the same time before meeting the HCl vapour.

Bromine diffusion demo. A gas jar of reddish-brown bromine is placed UNDER an empty gas jar of air, separated by a cover slip. When the cover slip is removed, the brown colour slowly spreads UPWARDS into the air jar until both jars are uniformly pale brown. This demonstrates random motion in all directions — bromine moves up despite being denser than air.

Dilution of coloured solutions. A drop of purple potassium manganate(VII) solution placed in a beaker of water slowly diffuses outward; eventually the whole beaker is a uniform pale purple. This is liquid diffusion — slower than gas diffusion because liquid particles move more slowly and have less room to manoeuvre.

Both evaporation and boiling turn a liquid into a gas, but the way it happens differs in three important ways:

Feature	Evaporation	Boiling
Where it happens	Only at the SURFACE	THROUGHOUT the bulk of the liquid
Temperature	At ANY temperature below the boiling point	Only at the FIXED boiling point
Bubbles?	No bubbles	Bubbles of vapour form inside the liquid

Why only the surface particles evaporate? Only at the surface can a particle that has enough kinetic energy actually escape into the gas phase. Particles below the surface that gain extra energy simply pass it on to neighbours via collisions before they can break free.

Cooling effect of evaporation. Because only the FASTEST-moving particles escape, the average kinetic energy of the remaining liquid DROPS — so the liquid COOLS. This is why sweating cools the body, why water evaporating off your skin feels cold, and why ethanol on the skin feels cold (lower boiling point → faster evaporation → larger cooling effect).

Boiling point is the temperature at which the vapour pressure of the liquid equals the atmospheric pressure. Above this temperature the entire liquid can flash to vapour and bubbles form throughout. Pure substances have a sharp, fixed boiling point; impurities raise the boiling point and broaden the range.