1. Chapter 5: GASES Part 2

2. Things to remember 1 atm = 760 mm Hg = 760 torr PV = nRT
𝑃 1 𝑉 1 𝑇 1 = 𝑃 2 𝑉 2 𝑇 2
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 + P

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. The Effusion of a Gas into an Evacuated Chamber
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).