States of Matter
Solids, Liquids, Gases, Plasma & Phase Changes
The Four States of Matter
Matter can exist in four distinct states — sometimes called phases. The state a substance is in depends on two factors: temperature and pressure. These factors determine how much energy the particles have and how close together they are.
In everyday chemistry you will mainly encounter solids, liquids, and gases. Plasma is the fourth state and the most common in the universe, though rarely encountered in a laboratory setting.
Particles are tightly packed in a regular, ordered arrangement. They vibrate in fixed positions but do not move past each other. Solids are incompressible and maintain their shape.
- High density; particles close together
- Very low compressibility
- No flow — rigid structure
- Examples: ice, iron, salt, diamond
Particles are close together but not in a fixed arrangement. They can slide past each other, giving liquids the ability to flow and conform to any container shape.
- High density; nearly as dense as solids
- Very low compressibility
- Flows freely; takes container shape
- Examples: water, mercury, ethanol, oil
Particles are far apart and move rapidly in random directions. Gases have no fixed shape or volume — they expand to fill any container. They are highly compressible.
- Very low density; particles far apart
- Highly compressible
- Fills any container completely
- Examples: oxygen, nitrogen, steam, CO₂
At extremely high temperatures, electrons are stripped from atoms, creating a soup of ions and free electrons. Plasma behaves differently from gas and responds to magnetic fields.
- Most common state in the universe
- Found in stars, lightning, auroras
- Used in neon signs, plasma TVs
- Requires temperatures of thousands of °C
Kinetic Molecular Theory
The Kinetic Molecular Theory (KMT) is the model that explains the behaviour of matter in terms of particles in constant motion. It was developed to explain the properties of gases, but its principles extend to solids and liquids.
All matter is made of tiny particles (atoms or molecules) that are in constant, random motion. The average kinetic energy of particles is directly proportional to temperature. Particles collide with each other and with container walls — these collisions are the origin of pressure in a gas.
| Property | Solid | Liquid | Gas |
|---|---|---|---|
| Particle arrangement | Regular, ordered lattice | Close, but disordered | Random, widely spaced |
| Particle motion | Vibrate in fixed positions | Slide and flow past each other | Move rapidly in all directions |
| Intermolecular forces | Very strong — hold structure rigid | Moderate — allow flow | Negligible — particles barely interact |
| Compressibility | Almost none | Almost none | High — large empty space between particles |
| Energy level | Lowest | Intermediate | Highest (of the three common states) |
Temperature is a measure of the average kinetic energy of particles in a substance. This is a crucial distinction: temperature measures the average, not the total. Two objects can have the same temperature but different total thermal energies if they have different masses or compositions.
Phase Changes
A phase change is the transition of matter from one state to another. Phase changes are physical changes — the chemical identity of the substance does not change, only the arrangement and energy of its particles.
All phase changes involve the absorption or release of energy. The temperature remains constant during a phase change, even as energy is being added or removed — all that energy is being used to break or form intermolecular forces, not to speed up particles.
Solid → Liquid. Occurs at the melting point. Requires energy input (endothermic) to overcome the rigid lattice structure.
Liquid → Solid. Occurs at the freezing point (same temperature as melting point). Energy is released as intermolecular bonds form.
Liquid → Gas. Can occur as evaporation (surface only, below boiling point) or boiling (throughout liquid, at boiling point).
Gas → Liquid. Energy is released as gas particles slow down and intermolecular forces pull them together. E.g. dew forming on glass.
Solid → Gas directly, without passing through the liquid state. E.g. dry ice (solid CO₂) sublimes at room temperature and pressure.
Gas → Solid directly. E.g. frost forming when water vapour deposits onto a cold surface without first becoming liquid.
Endothermic (absorb energy): melting, vaporisation, sublimation — going from lower to higher energy states.
Exothermic (release energy): freezing, condensation, deposition — going from higher to lower energy states.
The Heating Curve
A heating curve shows how the temperature of a substance changes as heat is continuously added. It reveals something important: temperature does not rise during phase changes. The energy added goes into breaking intermolecular forces instead of increasing kinetic energy.
Notice the two flat sections (plateaus) on the curve. These are the phase changes. During melting (at 0°C) and boiling (at 100°C for water at sea level), the temperature stays constant while all the added energy goes into overcoming intermolecular forces.
The energy required to complete a phase change without changing temperature is called latent heat. The latent heat of fusion applies to melting/freezing; the latent heat of vaporisation applies to boiling/condensation. Water has an unusually high latent heat of vaporisation — which is why sweating cools you so effectively.
Temperature & Pressure Effects
The state a substance adopts depends on both temperature and pressure. Increasing temperature gives particles more energy, pushing transitions toward the gas phase. Increasing pressure compresses particles, favouring the liquid or solid state.
Water boils at 100°C at sea level (1 atm pressure). At higher altitudes, atmospheric pressure is lower, so water boils at a lower temperature. On top of Mount Everest, water boils at around 70°C — which is why food takes longer to cook there.
For a pure substance, the melting point and the freezing point are the same temperature — the difference is only the direction of the phase change. Water melts at 0°C and freezes at 0°C. Impurities (like salt dissolved in water) lower the freezing point — the reason roads are salted in winter.
Worked Examples
Answer: During the melting phase change, the energy being added is not increasing the kinetic energy of the particles — instead, it is being used to break the intermolecular hydrogen bonds that hold the water molecules in the solid lattice structure. Until all those bonds are broken (i.e. all the ice has melted), the temperature cannot rise. This energy is called the latent heat of fusion.
Practice Questions
Q1. Which of the following correctly describes the particle arrangement in a liquid?
Q2. What happens to the temperature of water during the boiling phase change?
Q3. Dry ice (solid CO₂) transitions directly from solid to gas. This phase change is called:
Q4. Which of the following phase changes is exothermic (releases energy)?
Q5. Water boils at a lower temperature at high altitude because:
Key Takeaways
- Matter exists in four states: solid (fixed shape and volume), liquid (fixed volume, variable shape), gas (no fixed shape or volume), and plasma (ionised gas at extreme temperatures).
- The Kinetic Molecular Theory explains the properties of matter in terms of constantly moving particles whose average kinetic energy equals temperature.
- Phase changes are physical changes — the substance’s chemical identity remains the same, only the arrangement and energy of particles changes.
- The six phase changes: melting, freezing, vaporisation, condensation, sublimation, and deposition.
- Temperature remains constant during a phase change — all added energy goes into breaking or forming intermolecular forces (latent heat).
- A heating curve shows temperature vs heat added, with flat plateaus at the melting and boiling points.
- Boiling point decreases with lower atmospheric pressure — which is why water boils below 100°C at high altitude.