Transcript Slide 1
John E. McMurry • Robert C. Fay
C H E M I S T R Y
Chapter 9
Gases: Their Properties and Behavior
Lecture Notes
Alan D. Earhart Southeast Community College • Lincoln, NE
Stoichiometric Relationships with Gases 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 ?
2NaN 3 (s) 2Na(s) + 3N 2 (g)
Stoichiometric Relationships with Gases 2NaN 3 (s) 2Na(s) + 3N 2 (g)
Moles of N 2 produced:
45.0 g NaN 3 mol of N2
Volume of N 2 produced:
Use Ideal Gas Law
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?
Partial Pressure and Dalton’s Law 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.
Mole Fraction (X) = P total = P 1 + P 2 + … + P N Moles of component Total moles in mixture
n
i
X
i =
n
total or
P
i
X
i =
P
total
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
Example
A 1.00 L vessels contain 0.215 mole of N 2 mole of H 2 gas and 0.0118 gas at 25.5
o C. Determine the partial pressure of each component and the total pressure
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?
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.
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.
The Kinetic-Molecular Theory of Gases Copyright © 2008 Pearson Prentice Hall, Inc.
molar mass average speed Chapter 9/12
The Kinetic-Molecular Theory of Gases
Graham’s Law: Diffusion and Effusion of Gases Diffusion: The mixing of different gases by molecular motion with frequent molecular collisions.
Graham’s Law: Diffusion and Effusion of Gases Effusion: The escape of a gas through a pinhole into a vacuum without molecular collisions.
Graham’s Law:
Rate
a
1
m
Graham’s Law: Diffusion and Effusion of Gases
In comparing two gases at the same temperature and pressure
Rate 1 Rate 2
=
√m2 √m1
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
The Behavior of Real Gases
The volume of a real gas is larger than predicted by the ideal gas law.
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter 9/18
The Behavior of Real Gases
Attractive forces between particles become more important at higher pressures.
The Behavior of Real Gases
van der Waals equation
Correction for intermolecular attractions.
P
+
a n
2
V
2 V - n
b
= nRT Correction for molecular volume.
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