Kinetic Theory of Gases
Gases are made of tiny particles.
The particles are in constant motion.
They move in straight lines.
They collide with each other.
They also hit the walls of a container.
Key Ideas
- Particles are point masses.
- Their size is negligible.
- No forces act between them except during collisions.
- Collisions are perfectly elastic.
- Temperature is a measure of average kinetic energy.
Pressure
Particles strike the walls.
Each impact transfers momentum.
Many impacts per second create pressure.
Higher speed means higher pressure.
Temperature and Speed
Temperature rises with speed.
Average kinetic energy is ½ mv².
When temperature doubles, speed increases by √2.
Volume Effects
Particles spread out in a larger space.
Fewer collisions occur per unit area.
Pressure drops if volume increases at constant temperature.
Real Gases vs Ideal Gases
Real gases have finite size.
They attract each other weakly.
These factors matter at high pressure or low temperature.
Applications
The theory explains gas laws.
It guides engine design.
It helps predict weather patterns.
Summary
Gas particles move fast.
They bounce off each other.
Their motion creates pressure.
Temperature measures their energy.
The simple model works well for most gases.