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BEHAVIOR OF GASES
Unit 8
Chemistry
Langley
**Corresponds to Chapter 14 in the Prentice Hall Chemistry Book
PROPERTIES OF GASES
 No definite shape/volume
 Expands to fill its container
 Easily compressed (squeezed into a
smaller container)
 Compressibility is a measure of how much
the volume of matter decreases under
pressure
 Gases are easily compressed because of
the space between the particles in a gas
PROPERTIES OF A GAS
 Factors Affecting Gas Pressure
 Amount of Gas
 Increase amount, increase pressure
 Volume
 Reduce volume, increase pressure
 Temperature
 Increase temperature, increase pressure
 Relationship between pressure,
temperature, and volume is explained
through the Gas Laws
GAS LAWS
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Boyle’s Law
Charles’ Law
Gay-Lussac’s Law
Combined Gas Law
Ideal Gas Law
Dalton’s Law of Partial Pressure
Graham’s Law
BOYLE’S LAW
 If the temperature is constant, as the
pressure of a gas increases, the volume
decreases
 For a given mass of gas at constant
temperature, the volume of a gas varies
inversely with pressure
 As volume goes up, pressure goes down
 As volume goes down, pressure goes up
 P1V1 = P2V2
BOYLE’S LAW
 Real Life Example
 As you push on the end of a syringe, the
volume inside the syringe decreases as the
pressure on the syringe increases
 Mathematical Example 1:
 P1 = 758 torr
P2 = ?
V1 = 5.0L
V2 = 3.5L
BOYLE’S LAW
 Mathematical Example 2
 If 4.41 dm3 of nitrogen gas are collected at a
pressure of 94.2 kPa, what will the volume
be for this gas at standard pressure if the
temperature does not change?
CHARLES’ LAW
 As the temperature of an enclosed gas
increases, the volume increases, if the
pressure is constant
 The volume of a fixed mass of gas is directly
proportional to its Kelvin temperature if the
pressure is kept constant
 As volume goes up/down, temperature goes
up/down
 V1 = V2
T1
T2
Temperature must be in Kelvin!
CHARLES’ LAW
 Real Life Example
 Balloon Lab-As the temperature of the water
is increased, the volume of the balloon is
increased.
 Coke Can-Fill a coke can with a small
amount of water, as you heat the water
inside to near boiling, immediately invert the
coke can into ice-cold water so the coke can
is experiencing a dramatic drop in
temperature, volume of can will decrease
(can will crush in on itself)
CHARLES’ LAW
 Mathematical Example 1
 V1 = 250mL
V2 = 321mL
T1 = 300K
T2 = ?
 Mathematical Example 2
 With a constant pressure, the volume of a
gas is increased from 15.0L to 32.0L. If the
new temperature is 20.0°C, what was the
original temperature?
GAY-LUSSAC’S LAW
 As the temperature of an enclosed gas
increases, the pressure increases, if the
volume is constant
 The pressure of a gas is directly proportional to
the Kelvin temperature if the volume remains
constant
 P1 = P2
Temperature must be in Kelvin!
T1 T2
GAY-LUSSAC’S LAW
 Real Life Example
 Tires
 The faster a car goes, the higher the temperature
of the tire gets and the higher the pressure inside
the tires
 Mathematical Example 1
 P1 = ?
P2 = 789mmHg
T1 = 456K
T2 = 326K
GAY-LUSSAC’S LAW
 Mathematical Example 2
 The pressure in a tire is 1.8 atm at 20°C.
After a 200 mile trip, the pressure reading for
the tire is 1.9 atm. What is the temperature
inside the tire at that new pressure?
COMBINED GAS LAW
 Combines Boyle’s, Charles’, and Gay-Lussac’s
laws
 Describes the relationship among temperature,
pressure, and volume of an enclosed gas
 Allows you to perform calculation for situations
IF and ONLY IF the amount of gas is constant
 P1V1 = P2V2
Temperature must
be in
Kelvin!
T1
T2
IDEAL GAS LAW
 When you need to account for the number of
moles of gas in addition to pressure,
temperature, and volume, you will use the Ideal
Gas Equation
 Modified version of the Combined Gas Law
 PV = nRT
 n = number of moles
 R = ideal gas constant
 0.08206 (L-atm/mol-K)
 62.4 (L-mmHg/mol-K)
 8.314 (L-kPa/mol-K)
IDEAL GAS LAW
 Mathematical Example 1
 What is the pressure in atm exerted by 0.5
moles of N2 in a 10L container at 298
Kelvin?
 Mathematical Example 2
 What is the volume in liters of 0.250 moles
of O2 at 20°C and 0.974 atm?
IDEAL GAS LAW
 Mathematical Example 3
 What is the temperature of 76 grams of Cl2
in a 24L container at 890mmHg?
 Mathematical Example 4
 A deep underground cavern contains
2.24x106L of CH4 at a pressure of
1.50x103kPa and a temperature of 315K.
How many kilograms of CH4 does the cavern
contain?
IDEAL vs. REAL GASES
 Ideal gases follow the gas laws at all
conditions of pressure and temperature
 Conforms exactly to the all the assumptions
of the kinetic theory (no volume, no particle
attraction)doesn’t exist
 Real gases differ mostly from an ideal
gas at low temperature and high
pressure
 Under other conditions, behave as an ideal
gas would
DALTON’S LAW
 In a mixture of gases, the total pressure is the
sum of the partial pressure of the gases
 Partial pressure is the contribution each gas in a
mixture makes to the total pressure
 At constant volume and temperature, the total
pressure exerted by a mixture of gases is
equal to the sum of the partial pressures of the
component of gases
 Ptotal = P1 + P2 + P3 + …
DALTON’S LAW
 Mathematical Example 1
 In a container there are 4 gases with the
following pressures: Gas 1-2.5 atm, Gas 21.9 atm, Gas 3-798 mmHg, Gas 4-2.1 atm;
find the total pressure in the container.
DALTON’S LAW
 Mathematical Example 2
 In a sample of HCl gas, the pressure of the
gas is found to be 0.87 atm. If hydrogen
makes up 34% of the gas, what is the
pressure of the hydrogen?
GRAHAM’S LAW
 The ratio of the speeds of two gases at the
same temperature is equal to the square root
of the inverted molar masses
 The relative rate of diffusion
 Diffusion is the tendency of molecules to move
toward areas of lower concentration to areas of
higher concentration until the concentration is
uniform throughout
 Gases of lower molar mass diffuse and effuse faster
than gases of higher molar mass
 Effusion is when gas particles escape through tiny holes in
a container
GRAHAM’S LAW
 √(Molar MassB/Molar MassA)
 The rates of effusion of two gases are
inversely proportional to the square roots
of their molar masses
 Use periodic table to get molar masses
GRAHAM’S LAW
 Mathematical Example 1
 What is the ratio of the speeds of Helium
compared to Oxygen?
 Mathematical Example 2
 If Co2 has a speed of 22 m/s at 20°C, what is
the speed of HCl at the same temperature?