1. Gases

2. The Nature of Gases  1. Gases have mass –A car tire weighs more with air in it than it would completely empty.  2. It is easy to compress a gas. –If you apply pressure to a gas, the volume can be reduced considerably.  3. Gases fill their containers completely –Gas particles fill a container evenly

3. The Nature of Gases  4. Different gases can move through one another quite rapidly. –Diffusion  5. Gases exert pressure  6. The pressure of a gas depends on its temperature.

4. Kinetic Molecular Theory of Gases  1. Gases consist of small particles each with mass.  2. Size of the gas particle is insignificant compared to the distance between particles.  3. Gas particles are in constant, random motion

5. Kinetic Molecular Theory of Gases  4. Collisions between particles or between particles and a container are 100% elastic.  5. Kinetic energy of the gas particles depends only on the temperature of the gas.  6. Gas particles exert no force on one another. –No attraction between gas particles.

6. Measuring Gases: The 4 Variables  Amount of Gas (n): number of moles of gas  Volume (V): the area occupied by a gas. –A gas always completely fills its container so the volume of the gas equals the volume of the container.  Pressure (P): the force exerted by a gas on its container.  Temperature (T): measure in Kelvin

7. STP  Standard Temperature and Pressure –0˚C (273 K) –101.325 kPa (1 atm)  1 atmosphere of pressure is the average air pressure at sea level.

8. Boyle’s Law  Inverse relationship between pressure and volume. If the volume goes up, the pressure goes down. –Amount of gas and temperature are constant. P 1 V 1 = P 2 V 2

9. Boyle’s law  A sample of neon gas occupies a volume of 2.8 L at 1.8 atm. What will be its volume at 1.2 atm?

10. Boyle’s Law  What is the pressure of a gas if it was originally measured in a 7.8 Liter flask at 1 atm and is then moved to a 640 mL container?

11. Charles’ Law  There is a direct relationship between volume and temperature in gases. When temperature increases, so does the volume. –Pressure and amount of gas present must remain constant. V 1 /T 1 = V 2 /T 2

12. Charles’s Law  A gas is collected in a lab at 22˚C and occupies a volume of 18.6 L. The gas is cooled to 0˚C. What will be the final volume of the gas?

13. Charles’s Law  A bicycle tire was measured to have a volume of 400 cm 3 on a day when the temperature is 22˚ C. What will be the volume of the tire if the following day, the temperature is only 13˚C?

14. Combined Gas Law  Puts Boyle’s, Charles’, and Gay-Lussac’s Laws together. –Amount of gas must remain constant. P 1 V 1 /T 1 = P 2 V 2 /T 2

15. Homework  Page 913-914 –#1-7 and 9-11

16. The Ideal Gas Law  Combines all 4 variables of gases PV = nRT P= pressure V= volume N= # of moles of gas R= Universal gas constant T= Temperature (in Kelvin)

17. Dalton’s Law of Partial Pressures  The total pressure of a gas mixture is the sum of the pressures exerted by each individual gas. P total = P 1 + P 2 + P 3 + …  Need to subtract vapor pressure of water when a gas is collected over water or by water displacement

18.  Oxygen gas is collected at a temperature of 10˚C and a pressure of 1.02 atm. The volume of gas plus water vapor collected is 293 mL. How many moles of oxygen gas were collected? How many grams of oxygen gas were collected?

19. Gas Stoichiometry In a lab, we are mixing 50 grams of zinc with 103 grams of sulfuric acid. What volume of hydrogen can be produced from this reaction if the pressure is 1.13 atm and the temperature is 30 degrees Celsius?

20.  Hydrogen gas is produced from the reaction of magnesium with hydrochloric acid. 5.3 L of gas were collected by water displacement at a total pressure of 802 kPa and a temperature of 22 degrees Celsius. What mass of magnesium is required to react with excess acid in order to create this volume of gas.