1. 12.7 Laws and Models
Laws, such as the ideal gas law, predict how a gas will behave, but not why it behaves so.
A model (theory) explains why.

2. 12.8 The Kinetic Molecular Theory
Attempts to explain the behavior of an ideal gas.
Kinetic refers to the motion of the gas particles.
Molecular means the gases are composed of separate molecules (or atoms). 18

3. Postulates of the K-M Theory of Gases
Gases consist of tiny particles (atoms or molecules).
The particles are small compared to the average space between them. The volume of individual particles is negligible.
The particles are in constant random motion, colliding with the walls of their container. These collisions with the walls cause the pressure exerted by the gas.
The particles are assumed not to attract or repel each other.
The average kinetic energy of gas particles is directly proportional to the Kelvin temperature of the gas. 20

4. 12.9 Implications of the Kinetic Molecular Theory
The meaning of temperature:
The temperature of a gas reflects how rapidly its individual gas particles are moving.
At high temperatures the particles move very fast and hit the walls of their container more often. At low temperatures they move more slowly and collide with their container walls less often.
Temperature, then, is a measure of the motions of the gas particles.
The Kelvin temperature of a gas is directly proportional to the average kinetic energy of the gas particles. 21

6. The relationship between pressure and temperature:
As the temperature of a gas increases, the average speed of the molecules increases.
The molecules hit the sides of the container with more force (on average) and more frequently.
The net result is an increase in pressure.
Gay-Lussac’s Law 23

7. The relationship between volume and temperature
As the temperature increases the gas particles move faster, causing gas pressure to increase.
Assuming the gas is placed in a container with a moveable piston (fig ), the piston moves out to increase the volume of the container and keep the pressure constant.
Therefore, the volume of a gas will increase as temperature is raised at a constant pressure.
Agrees with experimental observations as summarized by Charles’ Law. 24

8. Gas Stoichiometry
This is similar to the stoichiometry problems in Ch 9 except that we now want to calculate the volume of a gas product as well as its mass.
Use the roadmap: grams  moles  moles  volume (L)
Example: 2KClO3(s)  2KCl(s) + 3O2(g)
If 10.5 grams of KClO3 decompose how many liters of O2 are produced?
Step 1: Convert grams of KClO3 to moles of KClO3
Step 2: Use the mole ratio to convert moles of KClO3 to moles of O2
Step 3: Use the ideal gas law to solve for the volume of O2 produced under the stated temperature and pressure.
𝑉= 𝑛𝑅𝑇 𝑃

9. Graham’s Law
Graham's Law deals with the effusion of gases.  This is not to be confused with diffusion which is the movement of molecules from a place of higher concentration to a place of lower concentration.
Effusion is the process in which a gas escapes through a small hole. The rate at which gases effuse (i.e., how many molecules pass through the hole per second) is dependent on their molecular weight (molar mass).
Effusion is random movement of gas molecules through a hole (or holes) in their container. A common example of this is a balloon filled with helium: first it is buoyant and floats in the air, but in a few days it hangs toward the ground or floats a few inches above the ground (if at all). The helium has escaped through the small holes in the balloon.