1. 15 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: 1515 where P 1, P 2 and P 3 represent the partial pressures of the gases in the mixture.

2. 16 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. 17 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. 18 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).

5. 19 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.

6. 20 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.

7. 21 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

8. 22 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. 12-13), 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.

9. Gas Stoichiometry 23

10. 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. 24