1. Lecture Notes Alan D. Earhart Southeast Community College Lincoln, NE Chapter 9 Gases: Their Properties and Behavior John E. McMurry Robert C. Fay CHEMISTRY

2. Stoichiometric Relationships with Gases 2Na(s) + 3N 2 (g)2NaN 3 (s) The reaction used in the deployment of automobile airbags is the high-temperature decomposition of sodium azide, NaN 3, to produce N 2 gas. How many liters of N 2 at 1.15 atm and 30.0 °C are produced by decomposition of 45.0 g NaN 3 ?

3. Stoichiometric Relationships with Gases 2Na(s) + 3N 2 (g)2NaN 3 (s) 45.0 g NaN 3  mol of N2 Volume of N 2 produced: Moles of N 2 produced: Use Ideal Gas Law

4. Examples Consider the reaction represented by the equation P 4 (s) + 6 H 2 (g)  4H 3 (g) What is the amount of P 4 is required to react with 5.39 L of hydrogen gas at 27.0 o C and 1.25 atm?

5. Partial Pressure and Dalton’s Law P total = P 1 + P 2 + … + P N Mole Fraction (X) = Dalton’s Law of Partial Pressures: The total pressure exerted by a mixture of gases in a container at constant V and T is equal to the sum of the pressures of each individual gas in the container. X i = P total PiPi X i = n total nini or Total moles in mixture Moles of component

6. Examples Determine the mole fractions and partial pressures of CO 2, CH 4, and He in a sample of gas that contains 0.250 mole of CO 2, 1.29 moles of CH 4, and 3.51 moles of He, and in which the total pressure is 5.78 atm

7. Example A 1.00 L vessels contain 0.215 mole of N 2 gas and 0.0118 mole of H 2 gas at 25.5 o C. Determine the partial pressure of each component and the total pressure

8. Example On a humid day in summer, the mole fraction of gaseous H 2 O (water vapor) in the air at 25.0 o C can be as high as 0.0287. Assuming a total pressure of 0.977 atm, what is the partial pressure (in atm) of H 2 O in the air?

9. The Kinetic-Molecular Theory of Gases 1.A gas consists of tiny particles, either atoms or molecules, moving about at random. 2.The volume of the particles themselves is negligible compared with the total volume of the gas; most of the volume of a gas is empty space. 3.The gas particles act independently of one another; there are no attractive or repulsive forces between particles.

10. The Kinetic-Molecular Theory of Gases 3.Collisions of the gas particles, either with other particles or with the walls of a container, are elastic (constant temperature). 4.The average kinetic energy of the gas particles is proportional to the Kelvin temperature of the sample.

12. The Kinetic-Molecular Theory of Gases Copyright © 2008 Pearson Prentice Hall, Inc. Chapter 9/12 average speed molar mass

13. The Kinetic-Molecular Theory of Gases

14. Graham’s Law: Diffusion and Effusion of Gases Diffusion: The mixing of different gases by molecular motion with frequent molecular collisions.

15. Graham’s Law: Diffusion and Effusion of Gases Graham’s Law: Effusion: The escape of a gas through a pinhole into a vacuum without molecular collisions.  Rate 1 m

16. Graham’s Law: Diffusion and Effusion of Gases In comparing two gases at the same temperature and pressure √m2 √m1 Rate 1 Rate 2 =

17. Example Determine how much faster Helium atoms moves, on average, than a carbon dioxide molecule at the same temperature Determine the molar mass and identity of a gas that moves 4.67 times as fast as CO 2

18. The Behavior of Real Gases Copyright © 2008 Pearson Prentice Hall, Inc. Chapter 9/18 The volume of a real gas is larger than predicted by the ideal gas law.

19. The Behavior of Real Gases Attractive forces between particles become more important at higher pressures.

20. The Behavior of Real Gases P V2V2 n2n2 = nRT+V - n van der Waals equation b a Correction for intermolecular attractions. Correction for molecular volume.

21. Example A sample of 3.50 moles of NH 3 gas occupies 5.20 L at 47 o C. Calculate the pressure of the gas (in atm) using A) the ideal gas equation B) the van der Waals equation a = 4.17 atm L/mol 2 b = 0.0371 L/mol