Gases
Physical Characteristics of Gases:
The Kinetic Theory (a model for gases):
1. Gases consist of a large number of tiny particles with
insignificant volume
2. The particles are in constant, random motion.
3. The collisions between particles and walls are elastic.
4. There are no forces of attraction or repulsion between
molecules
5. The average kinetic energy is directly proportional to
temperature in Kelvin ( __°C + 273).
Diffusion: spontaneous mixing due to random
motion (molecules moving from high
concentration towards low concentration)
Effusion: gas moving through a small hole
Real Gas – a gas that does not completely behave
according to the kinetic theory
Due to: 1. Molecules occupy space
2. exert attractive forces on each other
SOLID
LIQUID
GAS
Has its own
shape
Takes shape
of container
Fills
container
Highest
density
Middle
density
Lowest
density
Not
Not
Compressible
compressible compressible
Little
movement
Some
movement
Rapid
movement
Properties (P,V,T,n)
• Pressure (P): force that a gas exerts on a given
area
• 1 atm = 760 torr = 760 mmHg
• Volume (V): space occupied by gas
• 1 L = 1000 mL = 1000 cm3
• Temperature (T): measure of the average
kinetic energy of the gas
• MUST be in Kelvin!
• K = ˚C + 273
• Number of moles (n): quantity of gas
molecules
What is Pressure?
force Newtons
pressure 

area
cm 2
• Changing the force or area will change
the pressure (shoes!)
• Atmospheric (air) pressure is measured
by a barometer:
vacuum
mercury
atmospheric
pressure
• 1 atm = 760 torr
= 760 mmHg
=1.013 x 105 Pa
=101.3 kPa
• STP: standard temperature and pressure.
1 atm and 0ºC
Boyle’s Law
P1V1  P2V2
When T is constant: inverse
relationship between P and V (one
goes up… the other goes down)
A sample of oxygen gas occupies a
volume of 20 mL at 2.0 atm. At
what pressure will it occupy 55 mL?
P1V1
P2 
V2
2.0 atm  20mL
P2 
55 mL
P2  0.73 atm
Charles’s Law
V1 V2

T1 T2
When P is constant: direct
relationship between V and T (one
goes up… the other goes up)
V1 T2  T1 V2
You get a 1.7 L balloon inside at a
temperature of 23˚C. At what
temperature will the volume drop
to 1.5 L ?
Convert initial temperature to K
T1 = ˚C + 273
= 23 + 273
= 296 K
Rearrange equation
T1 T2
V1 V2
T1V2

 T2

V1
V1 V2
T1 T2
Solve for final temperature
296 K 1.5 L
 T2
1.7 L
T2  261 K
How cold is that?!?!
Gay-Lussac’s Law
P1 P2

T1 T2
P1T2  T1 P2
When V is constant: direct
relationship between P and T (one
goes up… the other goes up)
Higher T: more collisions in same
area.
A fire extinguisher has CO2 at 22ºC
and 20 atm. What is the pressure
at 30ºC?
T1  22  273  295K
P1  20 atm
T2  30  273  303K
P2  ?
P1 P2

T1 T2
P1T2
 P2
T1
P2  21atm
20 atm x 303 K
 P2
295 K
Practice!
• First 3 “Worksheets”
Combined Gas Law
P1V1 P2 V2

T1
T2
• True when moles are constant
• Use to remember
– Boyle’s law:
– Charles’s law:
– Gay-Lussac’s law:
• A balloon has a volume of 20.0 L at 23ºC
and 770 torr. What will its volume be at
685 torr and 25ºC?
P1  770torr T1  23  273  296K V1  20.0 L
P2  685torr T2  25  273  298K V2  ?
P1V1 P2 V2

T1
T2
V2 
P1T2 V1
 V2
P2 T1
(770 torr)(298 K)(20.0 L)
 23 L
(685 torr)(296 K)
Ideal Gas Law
PV=nRT
• R is the gas constant. It will
always have the same values.
• You must know which one to
use
• R = 8.314 kPa L K-1 mol-1 if
Pressure is in kilopascals(kPa),
Volume is in liters(L),
Temperature is in Kelvin(K)
R = 0.0821 L atm K-1 mol-1 if
Pressure is in
atmospheres(atm), Volume is
in litrers(L), Temperature is in
Kelvin(K)
Ideal Gas Law
• What volume is needed to store
0.050 moles of helium gas at 202.6
kPa and 400 K?
• What pressure will be exerted by
20.16 g hydrogen gas in a 7.5 L
cylinder at 20ºC?
• A 50 L cylinder is filled with argon
gas to a pressure of 10130.0 kPa at
30ºC. How many moles of argon
gas are in the cylinder?
• To what temperature does a 250
mL cylinder containing 0.40 g
helium gas need to be cooled in
order for the pressure to be 253.25
kPa?
Dalton’s Law of
Partial Pressure
PTotal  P1  P2  P3  ...
• The total pressure of a gas
sample is equal to the sum of
the partial pressures of
individual gases.
• Example: Earth’s atmosphere
Gas
Pressure (torr)
N2
593.5
O2
159.2
Ar
7.1
CO2
0.23
Total
760
93.4 kPa
3.3 kPa
Gases You Know
N2
The most common gas in our
atmosphere (78%)
Not reactive
O2
20% of the atmosphere
Supports combustion
CO2
Greenhouse gas
More dense than air
Used in fire extinguishers
H2
Very low density
Explosive if mixed with O2
P
T
V
Boyle’s Law
P
T
Charles’s Law
V
Vapor Pressure and
Boiling
• Vapor Pressure – the
pressure exerted by a vapor in
equilibrium with its liquid state.
• Liquid molecules at the surface
escape into the gas phase.
• These gas particles create
pressure above the liquid in a
closed container.
Vapor Pressure
Explained (Vid)
The condition in which two
opposing processes are occurring
simultaneously at equal rates is
called a dynamic equilibrium. A
liquid and its vapor are in
equilibrium when evaporation and
condensation occur at equal rates.
This can only be obtained in a
closed container.
• Vapor Pressure increases with
increasing temperature.
20oC
80oC
•As temperature increases, the
amount of vapor generated by a
liquid in a closed container
increases.
•This occurs because as the liquid
gains kinetic energy, the
molecules can overcome the
Evaporation vs Boiling
• Evaporation - when a liquid 
gas at any temperature
• Vaporization – When a liquid 
gas when heat is applied or at
the boiling temperature
• Boiling – occurs when the
vapor pressure above the liquid
equals the atmospheric
pressure.
Vapor Pressure Curves
Graph shows how boiling points change
with change in vapor pressure.
Boiling Points change with
pressure changes.
•Less pressure = lower boiling
point temperature (example =
water boils at lower