1. KINETIC MOLECULAR THEORY AND PRESSURE 13.3: pgs. 474 – 478 & 13.1: pgs. 442 - 445

2. Kinetic Molecular Theory  A model to explain the behavior of an ideal gas.  Composed of 5 Assumptions  Gases consist of tiny particles (atoms or molecules)  These particles are so small we say their size is zero  The particles are in constant random motion. When they hit the container, we get pressure.  Particles do not attract or repel one another.  Average kinetic energy of the gas particles is directly proportional to the Kelvin temperature of the gas

3. Temperature  Aussumption 5- Kelvin temperature is directly proportional to average kinetic energy of gas particles  Temperature is a measure of the movement of the gas particles At high temperatures, move fast, hit the walls more frequently How about low temps?? http://www.chm.davidson.edu/ChemistryApplets/KineticMolecularTheory/Maxwell.html

4. Relationship between Temperature and Pressure  The pressure of a gas is measured is due to collisions with the walls of a container  What happens to pressure as we heat a gas? (volume not allowed to change)  Pressure increase as temperature increases  How about when temp decreases?

5. Relationship between Volume and Temperature  Imagine a container that can change size  The pressure of the gas (P gas ) is equal to the pressure of the exterior (P ext ) surroundings.  What happens if we heat the gas to a higher temperature? http://www.youtube.com/watch?v=08ezkSrQ5lc

6. Real Gases  KMT is for ideal gases- those don’t exist  Ideal gas- particles have no attractions, zero volume  Why don’t these exist??  Real gases will behave like ideal gases under certain conditions  Low pressures – 1 atm or lower  Moderate temperatures – 0 o C or higher

7. Gases have Pressure  What causes gas pressure?  Pop can demo  It’s a measure of the collisions a gas has with a surface  Atmospheric pressure is measured in a barometer

8. Mercury Barometer  The pressure exerted by the atmospheric gasses on the surface of the mercury in the dish keep the mercury in the tube.  At sea level the height of the column of mercury averages 760 mm Hg  The level of Hg will change as the air pressure changes

9. Units of Pressure  Mercury barometers measure the height of the mercury column: mm Hg or torr  Standard atmosphere or atm  1 standard atmosphere = 1.000 atm = 760.0 mm Hg = 760.0 torr  SI unit: pascal, Pa  1 atm = 101,325 Pa  Engineering: pound per square inch, psi  1.000 atm = 14.69 psi  Can use these units to perform conversions

10. Atmospheric Pressure  Where does it come from?  Gravity pulling air down towards the earth  Influenced by weather conditions: lows and highs  Varies with altitude

11. DALTON’S LAW OF PARTIAL PRESSURES 13.1: Pgs. 445 - 457

12. Mixtures of Gases  Many important gases contain a mixture of components.  Examples Air Helium and Oxygen (Scuba divers’ tanks)  Studies of gaseous mixtures show that each component behaves independently of the others.

13. John Dalton  Known as Dalton’s Law of Partial Pressures  For a mixture of gases in a container, the total pressure exerted is the sum of the partial pressures of the gases present.  The partial pressure of a gas is the pressure that the gas would exert if it were alone in the container.  DOES NOT DEPEND ON THE TYPE of gas, just the NUMBER of particles present!!

14. Equation  For a mixture of 3 gases,  Subscripts refer to the individual gases (gas 1, gas 2, and gas 3) P total = P 1 + P 2 + P 3

15. Example  A gas mixture containing oxygen, nitrogen and carbon dioxide has a total pressure of 32.9 kPa. If PO 2 =6.6kPa and PN 2 = 23.0 kPa, what is PCO 2 ?