1. 12.6 Dalton’s Law of Partial Pressure
In mixtures of gases each component gas behaves independently of the other(s).
John Dalton (remember him from Ch 10) studied mixtures of gases. In 1803 he determined that, for a mixture of gases in a container, the total pressure exerted is the sum of the partial pressures of the individual gases present.
The partial pressure of a gas is the pressure the gas would exert if it were alone in the container. For example, in a mixture of three gases:
where P1, P2 and P3 represent the partial pressures of the gases
in the mixture.

2. The partial pressure of each gas can be calculated using the ideal gas law. Again, in an example with three gases:
The total pressure of a mixture of gases depends on the total number of moles of gas present, not the identity of the gases. This is because – as stated previously -- ideal gases behave the same.
See example on page 380.

3. 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.

4. 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

5. 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

6. 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

7. 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

8. 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

9. Gas Stoichiometry
This is similar to the stoichiometry problems in Ch 9, except that we want to calculate the volume of the gaseous product rather than the mass (in grams).
Therefore: grams  moles  moles  volume
Example: for the reaction
2KClO3(s)  2KCl(s) + 3O2(g)
If 10.5 grams of KClO3 decompose, how many liters of O2 are
produced?
Ans: Convert grams of KClO3 to moles of O2 (just like you did in Ch 9),
then use the ideal gas law to solve for volume (V) of O2 produced under
the given T and P conditions of the reaction. (See pages 385 – 386 in
text.)