Chapter 5: GASES Part 2
1
Things to remember
1 atm = 760 mm Hg = 760 torr
PV = nRT
𝑃1 𝑉1 𝑃2 𝑉2

=
𝑇1
𝑇2
STP conditions – 1 atm and 0°C
1 mol of a gas occupies a volume of 22.4
L at STP conditions.
2
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.
3
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.
4
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
5
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 + . . .
6
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
7
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
8
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
9
Kinetic Molecular Theory
Developed to explain gas behavior
1. Gases consist of a large number of molecules
in constant motion.
2. Volume of individual particles is  zero.
3. Collisions of particles with container walls
cause pressure exerted by gas.
4. Particles exert no forces on each other.
5. Average kinetic energy  Kelvin temperature
of a gas.
10
Kinetic Molecular Theory
As the kinetic energy increases, the
average velocity of the gas increases
11
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
12
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
13
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
14
Kinetic Molecular Theory:
App’s to Gases (continued)
The Effusion of a Gas into an Evacuated Chamber15
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.
16
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
17
Real Gases
18
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).
19