1. Kinetic Molecular Theory and Gases
KMT operates under 5 assumptions:
Gas particles are in constant, random, straight line motion.
Particles are separated by great distances.
Collisions are rapid and elastic.
No force between particles.
Total energy remains constant.
Natural laws are explained by theories. Gas law led to development of kinetic-molecular theory of gases in the mid-nineteenth century.
KMT Video
Average kinetic energy is directly proportional to temperature!! As temp goes up, so does kinetic energy

2. Kinetic-Molecular Theory
A gas consists of very small particles, each of which has a mass.
A gas spreads out and takes on the volume of its container. The volume of the gas particles is assumed to be zero because it is negligible compared with the total volume of the gas.
Gas particles are in constant, rapid, random motion. They move in straight lines, until they bump into something.
Temperature is a measurement of the average kinetic energy (speed) of the particles (Video)
Pressure is a measurement of the number and force of the collisions of the particles with the walls of the container
Gas particles do not lose energy like most other things in our world. When gases hit something they never slow down and stop.

3. Kinetic-Molecular Theory cont.
The collisions of gas particles with each other and with the container are totally elastic.
Gas particles exert no force on one another because their attractions are so weak they are assumed to be zero.

4. Gas Properties Relating to the Kinetic-Molecular Theory
Diffusion
Migration of molecules that results in a homogenous mixture.
Effusion
Escape of gas molecules through a tiny hole.

5. Graham’s Law
Molecules effuse through holes in a rubber balloon and deflate
This occurs at a rate equal to moles/time:
proportional to Temperature
inversely proportional to molar mass.
Therefore, He gas effuses more rapidly than O2 at the same T. He
Graham’s Law Demo

6. Graham’s Law Rate of effusion
“Proportional to”
Rate of effusion ְ
This also applies to diffusion, as heavier particles diffuse more slowly

7. Graham’s Law
Example:
Ammonia gas has a molar mass of 17.0 g/mol; hydrogen chloride gas has a molar mass of 36.5 g/mol. What is the ratio of their diffusion rates?

8. Graham’s Law
Example:
What does this mean? Ammonia will diffuse 1.47 times faster than HCl because the gas particles have a smaller mass

9. Graham’s Law Example on your own:
Carbon Monoxide gas is less massive than Carbon Dioxide gas. How much faster will CO diffuse compared to CO2?