Question

Why do gases exert pressure on the walls of their container?

Short answer

Gases exert pressure on the walls of their container because of the constant motion and collisions of gas particles with the container. These collisions are a result of the kinetic energy possessed by the gas particles. The pressure exerted by gases is also influenced by the volume of the container and the temperature of the gas.

Step-by-step solution

Understand the Kinetic Molecular Theory

The kinetic molecular theory is a model that describes the behaviour of gases. It is based on four postulates: 1) A gas consists of small particles (atoms or molecules) in constant random motion. 2) These particles are so small compared to the distances between them that the volume of individual particles can be considered negligible. 3) These particles are in constant collision with each other and the walls of the container. 4) The average kinetic energy of the particles is directly proportional to the temperature of the gas in Kelvin.

Connect the Theory to Gas Pressure

Gas pressure is a result of the gas molecules colliding with the walls of their container. Each collision exerts a small force on the wall. The constant motion and numerous collisions of these particles against the container wall cause the pressure that gases exert. The greater the number of collisions, the greater the pressure.

Understand the Influence of Volume and Temperature on Gas Pressure

From the kinetic theory of gases, two important factors influencing gas pressure are also identified. 1) Volume: If the volume of the container is decreased, gas particles have less space to move around, increasing the frequency of collisions and thus the pressure. 2) Temperature: As the temperature increases, the speed of the gas particles increases, leading to more frequent and forceful collisions, again increasing the pressure.

Key concepts

Gas Pressure

Gas pressure is the force exerted by gas particles when they collide with the walls of their container. Imagine countless tiny particles, like microscopic rubber balls, bouncing around inside a sealed box. Each time one of these particles hits the walls of the container, it applies a small force. Together, these countless tiny forces generate the overall pressure you can measure.

The concept is rooted in the Kinetic Molecular Theory, which explains the physical properties of gases through the motion of particles. Since gas molecules are in constant random motion, they collide incessantly with the container walls, translating this activity into pressure. It's a little bit like a very lively crowd bumping against the fences at a concert—each bump adds up, impacting the fence, or in this case, the container walls.

This is why, even though we can't see the particles, we can feel or measure the pressure they exert. It's this force from the countless collisions that explains how tires stay inflated, balloons stay rounded, and why changes in conditions can influence pressure.

Collision with Container Walls

Collisions between gas particles and container walls are crucial to understanding gas behavior. These collisions are perfectly elastic, meaning no energy is lost; it just keeps the particles bouncing around energetically.

When a gas particle collides with the container wall, it rebounds, exerting an outward force on the wall. The result of many particles doing this millions of times each second is gas pressure.

The energy and momentum of particles ensure that they continuously impact and rebounce, maintaining consistent pressure within the container. The more frequent and forceful these collisions are, the greater the gas pressure becomes.

• More collisions mean more pressure.
• Energy is conserved and transferred through these particle impacts.

Influence of Volume on Gas Pressure

Volume plays a pivotal role in determining gas pressure. According to the kinetic molecular theory, gas particles move freely and fill the volume of their container.

If the volume of the container decreases—imagine a balloon being squeezed—there's less space for the particles to move. In this cramped environment, particles are more likely to collide with the walls and each other more frequently, thus increasing the gas pressure.

This relationship is described by Boyle's Law, which states that at a constant temperature, the pressure of a gas is inversely proportional to its volume:

• Decrease in volume → Increase in pressure
• Increase in volume → Decrease in pressure

Understanding how changes in volume affect gas pressure is essential for various real-world applications, from inflating tires to working with pressurized systems in laboratories and factories.

Influence of Temperature on Gas Pressure

Temperature directly affects the kinetic energy of gas particles. According to the Kinetic Molecular Theory, the average kinetic energy of gas particles is proportional to the temperature measured in Kelvin.

When temperature increases, particles gain more kinetic energy and move faster. This increase in speed results in more frequent and more energetic collisions with the container walls, thereby raising the pressure.

This relationship between temperature and pressure is encapsulated in Gay-Lussac's Law:

• Increase in temperature → Increase in pressure
• Decrease in temperature → Decrease in pressure

Practical implications of this include understanding why a tire might burst on a hot day or why a sealed container might explode if heated excessively. These phenomena highlight how critical it is to monitor temperature changes in contexts where gas pressure is a factor.