3.2. Kinetic Theory & Gas Laws (B3)

Macroscopic properties like pressure, temperature, and volume arise from the collective behavior of huge numbers of gas molecules:

• Pressure: Caused by collisions of gas molecules with the walls of a container. From kinetic theory: pressure relates to average molecular momentum change.

• Temperature: Directly linked to the average kinetic energy of the molecules.

• Volume: Reflects the space occupied by molecules moving freely.

• Number of molecules N is related to amount of substance n using Avogadro’s number

The ideal gas laws come from observations like:

• • At constant temperature: pV=constant

  • At constant pressure: V∝T (Charles' Law)

  • At constant volume: p∝T (Pressure Law)

• Assumptions of the kinetic theory of gases:

  • Gas molecules are small compared to the distances between them.

  • No forces act between molecules except during brief elastic collisions.

  • Collisions are perfectly elastic (no energy loss).

  • Molecules are in random continuous motion.

  • The volume of the gas molecules themselves is negligible compared to the container volume.

These lead to the Ideal Gas Law: pV=nRT

• Random, erratic movement of small particles (like pollen grains) suspended in a fluid. Observed by Robert Brown, explained by collisions with invisible gas/liquid molecules. Supports the idea that matter is made of tiny particles in constant motion — a foundation of kinetic theory.

• When gas volume is decreased at constant temperature, pressure increases. Observed in experiments using sealed syringes or pressure sensors. Supports that gas molecules move randomly and that pressure arises from collisions with container walls.

• Heating a fixed-volume gas causes its pressure to rise. Observed with pressure gauges as gas is heated in a rigid container. Supports the idea that temperature reflects average kinetic energy of molecules.