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, NaN3, to produce N2 gas. How many liters of N2 at
1.15 atm and 30.0 °C are produced by decomposition of
45.0 g NaN3?
2NaN3(s)
2Na(s) + 3N2(g)
Stoichiometric Relationships
with Gases
2NaN3(s)
Moles of N2 produced:
45.0 g NaN3

mol of N2
Volume of N2 produced:
Use Ideal Gas Law
2Na(s) + 3N2(g)
Examples
 Consider the reaction represented by the equation
P4(s) + 6 H2(g)  4H3(g)
What is the amount of P4 is required to react with 5.39 L of
hydrogen gas at 27.0oC 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.
Ptotal = P1 + P2 + … + PN
Moles of component
Mole Fraction (X) =
Total moles in mixture
Xi =
ni
ntotal
or
Xi =
Pi
Ptotal
Examples
 Determine the mole fractions and partial pressures of CO2,
CH4, and He in a sample of gas that contains 0.250 mole
of CO2, 1.29 moles of CH4, 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 N2 gas and 0.0118
mole of H2 gas at 25.5oC. Determine the partial pressure
of each component and the total pressure
Example
 On a humid day in summer, the mole fraction of gaseous
H2O (water vapor) in the air at 25.0oC can be as high as
0.0287. Assuming a total pressure of 0.977 atm, what is
the partial pressure (in atm) of H2O 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
molar
mass
average
speed
Copyright © 2008 Pearson Prentice Hall, Inc.
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 CO2
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 n2
V-n
b = nRT
V2
Correction for
molecular volume.
Example
 A sample of 3.50 moles of NH3 gas occupies 5.20 L at
47oC. 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/mol2
b = 0.0371 L/mol