State the distinguishing properties of solids, liquids and gases

IGCSE Chemistry 0620 – States of Matter

Learning objective

State the distinguishing properties of solids, liquids and gases and use the particle‑theory model to explain their behaviour, changes of state, the effect of temperature & pressure on gases, and diffusion.

1. Distinguishing properties (syllabus core)

For each state the following properties are listed exactly as required by the syllabus.

• Shape – Solids have a definite shape; liquids and gases take the shape of their container.
• Volume – Solids and liquids have a definite volume; gases expand to fill the container.
• Particle arrangement – Solids: particles are closely packed in a lattice (crystalline) or in a disordered arrangement (amorphous).
  Liquids: particles are close together but not in fixed positions.
  Gases: particles are widely spaced and move independently.
• Compressibility – Solids are very incompressible, liquids are slightly compressible, gases are highly compressible.
• Diffusion rate – Very slow in solids, moderate in liquids, very fast in gases.

2. Particle‑theory description of each state

The particle‑theory model links the macroscopic properties above to what is happening at the molecular level.

• Solids
  • Particles are tightly packed in a regular lattice (crystalline) or an irregular arrangement (amorphous).
  • Average separation is the smallest of the three states.
  • Each particle vibrates about a fixed position; translation and rotation are negligible.
  • Strong intermolecular forces keep the particles locked in place.
• Liquids
  • Particles are close together but not fixed; they can slide past one another.
  • Average separation is slightly larger than in solids.
  • Particles vibrate, rotate and translate (slide).
  • Intermolecular forces are moderate – enough to give a definite volume but weak enough to allow flow.
• Gases
  • Particles are far apart and move independently.
  • Average separation is greatest; the volume occupied depends entirely on the container.
  • Rapid, random translational motion; collisions with each other and the walls produce pressure.
  • Intermolecular forces are negligible.

3. Comparison of the three states

Property	Solid (s)	Liquid (l)	Gas (g)
Shape	Definite shape	Takes shape of container	Takes shape of container
Volume	Definite volume	Definite volume	Expands to fill container
Particle arrangement	Closely packed in a lattice (crystalline or amorphous)	Close together, no fixed positions	Widely spaced, free movement
Particle motion	Vibration only	Vibration, rotation, translation (sliding)	Rapid random translation
Compressibility	Very low	Low	High
Diffusion rate	Very slow	Moderate	Very fast
Examples	Ice, iron, sugar	Water, ethanol, mercury	Steam, O₂, CO₂

4. Changes of state (core requirement)

• Melting (s → l) – particles gain kinetic energy, overcome part of the intermolecular forces and begin to slide.
• Freezing (l → s) – loss of kinetic energy allows forces to lock particles into fixed positions.
• Boiling (l → g, bulk) – particles acquire enough energy to break all attractive forces throughout the liquid.
• Condensation (g → l) – cooling reduces kinetic energy so particles can come close enough for intermolecular forces to act.
• Evaporation (l → g, surface) – only the most energetic surface particles escape at temperatures below the boiling point.

Suggested diagrams: energy‑level diagram for each change and a heating‑cooling curve showing the solid, liquid and gas regions with labelled melting point (M) and boiling point (B). Use the state symbols s, l, g in all equations.

5. Effect of temperature & pressure on gas volume (core requirement)

The three individual gas laws required by the syllabus are listed below, followed by a short worked example.

Worked example (Gay‑Lussac’s law):

A sealed 2.0 L container holds nitrogen at 1.0 atm and 300 K. The temperature is raised to 450 K while the volume remains fixed. Calculate the new pressure.

\[ \frac{P_1}{T_1} = \frac{P_2}{T_2} \;\;\Rightarrow\;\; P_2 = P_1 \times \frac{T_2}{T_1} = 1.0\;\text{atm} \times \frac{450}{300} = 1.5\;\text{atm} \]

6. Diffusion (core requirement)

• Definition – net movement of particles from an area of higher concentration to one of lower concentration.
• Why it occurs – particles are in constant random motion (kinetic theory) and collide, spreading them through space.
• Effect of molecular mass – at a given temperature lighter molecules move faster, so diffusion rate \(\propto 1/\sqrt{M}\) where \(M\) is the molar mass.

Simple calculation (compare hydrogen and carbon‑dioxide):

\[ \frac{\text{rate}_{\mathrm{H_2}}}{\text{rate}_{\mathrm{CO_2}}} = \sqrt{\frac{M_{\mathrm{CO_2}}}{M_{\mathrm{H_2}}}} = \sqrt{\frac{44}{2}} \approx 4.7 \] Thus hydrogen diffuses about five times faster than carbon‑dioxide under the same conditions.

7. Beyond the core (optional – for extended learners)

Kinetic‑particle‑theory explanation of phase changes – heating adds kinetic energy; when the average kinetic energy equals the energy required to overcome intermolecular forces, a phase change occurs (e.g., melting point, boiling point). Energy‑level diagrams can illustrate the latent‑heat plateaus on a heating‑cooling curve.

Derivation of the ideal‑gas equation – starting from the kinetic theory, pressure is derived as \(P = \frac{1}{3}\frac{N}{V}mv_{\text{rms}}^{2}\). Substituting \( \frac{1}{2}mv_{\text{rms}}^{2} = \frac{3}{2}k_{\mathrm B}T\) leads to \(PV = nRT\).

8. Summary

Solids, liquids and gases differ in shape, volume, particle arrangement, compressibility and diffusion. These differences arise from the balance between kinetic energy and intermolecular forces, which the particle‑theory model describes. Understanding this model enables us to explain phase changes, the quantitative relationships governing gases (Boyle’s, Charles’s, Gay‑Lussac’s laws and the ideal‑gas equation), and why diffusion rates depend on molecular mass.