The pressure exerted by a gas is a macroscopic property that originates from the relentless, chaotic motion of its microscopic constituent particles.1 According to the kinetic particle model, a gas is composed of a vast number of particles in constant, random motion.2 The pressure we measure is the cumulative result of these particles colliding with the walls of their container.4 Each collision imparts a small force, and the sum of these forces over the container's surface area creates a steady pressure.7 Understanding this particle-level mechanism is key to explaining how changes in a gas's conditions, such as temperature and volume, affect its pressure.

When a fixed mass of gas is confined to a container of constant volume, its pressure is directly proportional to its absolute temperature.8 This relationship, known as the Pressure Law (or Amontons's Law), can be explained by examining the behavior of the gas particles as they are heated or cooled.10

Increasing the Temperature

When the temperature of the gas is increased, thermal energy is transferred to its particles, increasing their average kinetic energy.5 The key consequences of this are:

• Increased Particle Speed: An increase in kinetic energy means the gas particles move faster, on average.5
• More Forceful Collisions: Because the particles are moving faster, they have greater momentum. When they strike the container walls, the change in momentum during the collision is larger, resulting in a greater force being exerted on the wall for each individual impact.13
• More Frequent Collisions: Faster-moving particles will travel across the container in less time, leading to them colliding with the walls more often.5

The combined result of collisions being both more forceful and more frequent is an increase in the total average force exerted on the container walls. Since pressure is defined as force per unit area, and the area of the container is constant, the gas pressure increases.3

Decreasing the Temperature

Conversely, if the temperature of the gas is decreased, the particles lose kinetic energy and slow down.5 This leads to:

• Collisions with the walls that are less forceful.
• Collisions with the walls that are less frequent.

This reduction in the rate and force of collisions results in a lower total force on the walls, causing the pressure to decrease.5

When the volume of a fixed mass of gas is changed while its temperature is held constant, its pressure changes in inverse proportion. This relationship is known as Boyle's Law.3 The explanation again lies in the frequency of particle collisions.

Decreasing the Volume

If the volume of the container is decreased (the gas is compressed), the same number of particles are now confined to a smaller space.13 While the temperature remains constant, meaning the average speed of the particles and the force of each individual collision do not change, the following occurs:

• Increased Collision Frequency: The particles have less distance to travel before striking a wall. This significantly increases the rate, or frequency, at which the particles collide with the container walls.13

Because the walls are being struck more often, the total average force exerted on the walls increases. This increased force is acting on a smaller surface area, both of which contribute to a rise in pressure.3

Increasing the Volume

If the volume of the container is increased, the particles have more space to move around.3

• Decreased Collision Frequency: The particles must travel further to reach a wall, which means they collide with the walls less frequently.3

Since the temperature and the number of particles are unchanged, the average force of each collision remains the same. However, with fewer collisions occurring over a given time, the total average force on the walls decreases, resulting in a lower gas pressure.3