1. Why Assumptions Are Needed
The kinetic theory of gases is built on a set of simple assumptions about gas molecules. These assumptions help explain gas behaviour using microscopic ideas. Although real gases do not follow all of these perfectly, they work very well for ideal gases.
2. Gas Molecules Are Extremely Small
The size of each molecule is so tiny compared to the distance between them that we treat molecules as point masses. This means their actual volume is negligible compared to the total volume of the gas.
2.1. Why This Assumption Helps
It allows us to ignore the actual size of particles when calculating pressure, volume, and molecular speeds.
3. Large Intermolecular Distances
The distance between molecules is much larger than the size of the molecules themselves. This creates large empty spaces inside gases, explaining their high compressibility.
3.1. Example
Air inside a syringe can be compressed easily because molecules can be pushed closer together into the empty spaces.
4. No Intermolecular Forces
Gas molecules are assumed to have no attraction or repulsion toward each other except during collisions. They move independently unless they collide.
4.1. Implication
This explains why gases expand to fill any available space and why they do not stick together like liquids or solids.
5. Molecules Move in Random, Straight-Line Motion
Between collisions, gas molecules move in straight lines at constant speed. Their directions change only when they collide with another molecule or a container wall.
5.1. Reason
Since there are no forces acting between collisions, the motion remains uniform until disturbed.
6. Collisions Are Perfectly Elastic
When gas molecules collide with each other or with the container walls, there is no loss of kinetic energy. The total kinetic energy before and after the collision remains the same.
6.1. Important Note
Although individual molecules may gain or lose energy in a collision, the total energy of all molecules remains constant.
7. Duration of Collisions Is Negligible
Collisions happen extremely quickly, so the time spent colliding is negligible compared to the time spent moving freely between collisions.
8. Average Kinetic Energy Depends Only on Temperature
The average kinetic energy of gas molecules is directly proportional to the absolute temperature (in Kelvin). This means:
\( KE_{avg} \propto T \)
Higher temperature → faster molecular motion.
8.1. Meaning
Temperature is basically a measure of the average kinetic energy of gas molecules.
9. Number of Molecules Is Extremely Large
A gas contains an enormous number of molecules. This allows us to use statistical methods to predict average behaviour, even though individual molecules behave randomly.
9.1. Example
A small volume of air contains more molecules than the number of stars in the observable universe. With such huge numbers, averages become reliable.