Chapter 5: GASES Part 2

1. Chapter 5: GASES Part 2

2. Things to remember • 1 atm = 760 mm Hg = 760 torr • PV = nRT • = • STP conditions – 1 atm and 0°C • 1 mol of a gas occupies a volume of 22.4 L at STP conditions.

3. Gas Stoichiometry • Example 5.12 CaO is produced by the thermal decomposition of calcium carbonate. Calculate the volume of CO2 at STP produced when 152 g of CaCO3 decomposes.

4. Limiting Reagent Stoich • 5.13 A sample of methane gas (CH4) having a volume of 2.80 L at 25° C and 1.65 atm was mixed with a sample of oxygen gas having a volume of 35.0 L at 31° C and 1.25 atm. The mixture was then ignited to form carbon dioxide and water. Calculate the volume of CO2 formed at a pressure of 2.50 atm and a temperature of 125° C.

5. Molar Mass of a Gas • n = grams of gas / molar mass • P = nRT/ V = (grams)RT / (molar mass)V AND • d = mass / V SO • P = dRT/ molar mass • Rearranged: Molar mass = dRT / P

6. Dalton’s Law of Partial Pressures • Since gas molecules are so far apart, we can assume that they behave independently. • Dalton’s Law: in a gas mixture, the total pressure is the sum of the partial pressures of each component: PTotal = P1 + P2 + P3 + . . .

7. Using Dalton’s Law: Collecting Gases over Water • Commonly we synthesize gas and collect it by displacing water, i.e. bubbling gas into an inverted container

8. Using Dalton’s Law: Collecting Gases over Water • To calculate the amount of gas produced, we need to correct for the partial pressure of water: Ptotal = Pgas + Pwater

9. Using Dalton’s Law: Collecting Gases over Water Example 3: Mixtures of helium and oxygen are used in scuba diving tanks to help prevent “the bends”. For a particular dive, 46 L of He at 25°C and 1.0 atm and 12 L of O2 at 25°C and 1.0 atm were each pumped into a tank with a volume of 5.0 L. Calculate the partial pressure of each gas and the total pressure in the tank at 25°C

10. Kinetic Molecular Theory Developed to explain gas behavior • Gases consist of a large number of molecules in constant motion. • Volume of individual particles is  zero. • Collisions of particles with container walls cause pressure exerted by gas. • Particles exert no forces on each other. • Average kinetic energy  Kelvin temperature of a gas.

11. Kinetic Molecular Theory • As the kinetic energy increases, the average velocity of the gas increases

12. Kinetic Molecular Theory: Applications to Gases • As volume of a gas increases: • the KEavg of the gas remains constant. • the gas molecules have to travel further to hit the walls of the container. • the pressure decreases

13. Kinetic Molecular Theory: App’s to Gases (continued) • If the temperature increases at constant V: • the KEavg of the gas increases • there are more collisions with the container walls • the pressure increases

14. Kinetic Molecular Theory: App’s to Gases (continued) • effusion is the escape of a gas through a tiny hole (air escaping through a latex balloon) • the rate of effusion can be quantified

15. Kinetic Molecular Theory: App’s to Gases (continued) The Effusion of a Gas into an Evacuated Chamber

16. Kinetic Molecular Theory: App’s to Gases (continued) • Diffusion: describes the mixing of gases. The rate of diffusion is the rate of gas mixing. • Diffusion is slowed by gas molecules colliding with each other.

17. Real Gases Real Gases do not behave exactly as Ideal Gases. For one mole of a real gas,PV/RT differs from 1 mole. The higher the pressure, the greater the deviation from ideal behavior

18. Real Gases

19. Real Gases • The assumptions of the kinetic molecular theory show where real gases fail to behave like ideal gases: • The molecules of gas each take up space • The molecules of gas do attract each other • Chemists must correct for non-ideal gas behavior when at high pressure(smaller volume) and low temperature(attractive forces become important).