Standard 4: Gases.
Chemistry.
Ms. Siddall.
vocabulary
1. Temperature
2. Heat
3. Absolute zero
4. Kelvin scale
5. Standard Temperature
and Pressure
6. Atmospheric pressure
7. Volume
8. Density
8. Kinetic theory
9. Kinetic energy
10.Diffusion
11.Effusion
12.Homogeneous
13.Dalton’s law of
partial pressure
14.Combined gas law
Standard 4f: Lowest Possible Temperature
Absolute zero
= temperature where there is no
particle movement (no kinetic energy).
There can be no lower temperature.
absolute zero = 0 Kelvin = 0K
Deep space ~ 2.7K
Summary 1
At what temperature do particles stop
moving?
Standard 4e: Kelvin/Celsius conversions
Kelvin temperature = °C + 273
 Example: room temperature is about 22°C.
What is that in Kelvin?
Kelvin = °C + 273 = 22 + 273
= 295K
 Example: What is 0K in °C?
°C = Kelvin - 273 = 0 - 273
= -273°C
Summary 2
1. Convert 100°C to Kelvin
2. Convert 100K to Celsius
Standard 4d:
Standard Temperature and Pressure
STP:
1 mole of gas = 22.4L
Temperature = 0°C = 273K
Pressure = 1 atm
Summary 3
 One mole of gas at 0°C has a volume
of 22.4L. What is the pressure of the
gas?
4a: Pressure
Atmospheric pressure = force of air
hitting you
 Particles are in constant motion
 Particles have mass
 Mass x motion = force
 Force = pressure
Summary 4
 What causes atmospheric pressure?
Measuring pressure
 Atmospheric pressure = 1atm (atmosphere)
 1atm = 760 Torr
= 760mmHg
= 29.9inHg
= 101,325 Pa
= 101.325kPa
= 14.7psi (lb/in2)
 Example: Coca-Cola© is bottled at approximately
2 atmospheres. Convert to Torr.
2 atm
760 Torr
1 atm
~ 1520
Torr
Summary 5
Convert 2atm to:
1. Psi
2. kpa
Liquids and gases compared
• Atmospheric pressure = 14.7psi
• Deep ocean pressure = 16,000psi
• Water molecules are close together
(high pressure) and air particles are far
apart (lower pressure)
• Density of water = 1g/ml = 1000g/L
• Density of air = 1g/L
Summary 6
 Explain the difference between
molecules in a liquid and molecules in
a gas.
Kinetic Molecular Theory:
1. Gases consist of small particles that have
mass.
2. The volume of a gas is mostly empty space.
3. There are no intermolecular attractions
between gas particles.
4. Particles are in constant, random, rapid,
straight line motion.
5. Collisions are elastic (no loss of energy).
6. Kinetic energy of particles depends on
temperature.
Summary 7
 Which two points of the Kinetic Molecular
Theory explain pressure?
(remember pressure = mass x motion)
Temperature measures average kinetic energy.
Individual particles have different energies.
Most molecules have
average kinetic energy
Summary 8
 According to the graph: Which temperature
shows the highest average kinetic energy?
Summary 9
Which points of the Kinetic Molecular Theory
explain these properties?
 Gases are compressed easily
 Gases fill their containers
4b: Diffusion
Diffusion = ability of gases to mix
 Particles move constantly
 Particles are not attracted to each other
Therefore:
 Particles fill a container
 Particles form a homogeneous mixture


Lighter gases diffuse more quickly
Heavy gases diffuse more slowly
Summary 10
 What is a homogeneous mixture?
Summary 11
 Helium gas and oxygen gas are released
into a room at the same time. Which gas fills
the room first? Explain why.
Dalton’s Law
 The total pressure of a mixture of gases is
the sum of the pressure of each gas:
 Ptotal = P1 + P2 + P3 + ….
Summary 12
At the summit of mount Everest the total
pressure is about 0.3atm.
 Assume PTotal = PN2 + PO2
 If PN2 = 0.25atm what is the pressure of
oxygen?
4c: The Gas Laws
The combined gas law:
for changing conditions
P1V1 = P2V2
T1
T2
Initial conditions
P1
V1
Initial pressure (atm)
T1
Initial temperature* (K)
Initial volume (L)
Final conditions
P2
V2
Final pressure (atm)
T2
Final temperature* (K)
Final volume (L)
*Temperatures must be in Kelvin only
 Example: a gas at STP has a volume of 10L. The
pressure is increased to 5atm. Temperature
remains constant. What is the new volume?
Initial conditions
P1 1 atm
Final conditions
P2 5 atm
V1 10L
V2
?
T1 273K
T2
= T1
P1V1 = P2V2
T1
T2
P1V1T2 = P2V2T2
T1P2
T2P2
 Example: A plane flies at 35,000 feet. Pressure
outside is 0.25atm. Temperature in the cabin is
27°C. What will the temperature be if the cabin loses
pressure?
Initial conditions
Final conditions
P1 1 atm
P2
0.25 atm
V1
=V2
V2
=V1
T1
300K
T2
?
P1V1 = P2V2
T1
T2
T1
= T2
P1V1
P2V2
Summary 13
 A gas has a fixed volume and a pressure of
1atm. If temperature increases from 200K to
400K what is the new pressure?
Relationships using the Gas Law.
Pressure vs. Volume
VP
VP
Pressure (atm)
= inverse relationship = inversely proportional
Smaller volume
= more collisions
= more pressure
volume (L)
Summary 14
If the volume is decreased by a factor of 3,
what happens to the pressure?
Temperature vs. pressure
pressure (atm)
T  P 
T  P 
= direct relationship = directly proportional
Higher temperature
= more energy
= more collisions
= more pressure
Temperature (K)
Summary 15
If the temperature is tripled, what happens to
the pressure?
Volume vs. Temperature
Volume (L)
T  V 
T  V 
= direct relationship = directly proportional
Temperature (K)
Higher temperature
= more energy
Gas takes up more
space
Summary 16
If the temperature is decreased by a factor of
4 what happens to the volume of the gas?
moles vs. pressure
 n  (=number of moles) P 
n P
= direct relationship
= directly proportional
More gas = more collisions = more pressure
Summary 17
If the number of particles is doubled, what
happens to the pressure?
Honors Only
4h: solve problems by using the ideal
gas law in the form PV = nRT
Ideal Gas Law
PV = nRT




P = pressure (atm)
V = volume (L)
n = number of moles
R = universal gas constant
= 0.0821L·atm/mol·K
 T = temperature (K)
 Example: what is the volume of 1 mole of
gas at standard temperature and pressure?
 PV = nRT
 V = nRT/P
 V = (1 mol)(0.0821L·atm) (273.15K)
(mol·K) (1 atm)
 V = 22.4L
Summary 18
 At 500°C what is the pressure of one mole
of gas that has a volume of 10L?
The Ideal Gas Law (and the Combined Gas Law) only
work under ‘ideal conditions’
Ideal conditions = moderate temperature & pressure
Non-ideal conditions: (when the gas laws no longer
makes good predictions)
 Extremely high pressure:
 Particles are too close together
 Particles take up space
 Extremely low temperature:
 Particles have very low energy
 Particles are attracted to each other
Summary 19
Explain why the inverse relationship between
pressure and volume is not linear
4i:apply Dalton’s Law of Partial Pressures to
describe the composition of gases and
Graham’s Law to predict diffusion of gases
 Graham’s Law: rate = (1/mass)
 The rate of diffusion depends on particle mass
 Gases with lower mass diffuse faster
 Effusion: rate of gas passing through a small
opening in a container
Summary 20
 Which gas will effuse faster; chlorine or
bromine? Why?
4g: the kinetic molecular theory relates
the absolute temperature of a gas to the
average kinetic energy of its molecules
or atoms
 At 0K all motion ceases
 At 0K: average kinetic energy = zero
Summary 21
 If the average kinetic energy of a gas is
zero, what is the temperature of that gas in
Kelvin? Why?