Gases Notes
12.1 A. Physical Properties:
1. Gases have mass. The density is much
smaller than solids or liquids, but they
have mass. (A full balloon weighs more
than an empty one.)
2. Gases can be compressed. It is very easy
to reduce the volume of a gas.
Example: air in a bicycle tire
3. Unlike liquids, gases completely
fill their containers.
4. Gases can move through each other
rapidly. Example: the diffusion of food
smells and perfume.
5. Gases exert pressure.
6. The pressure of a gas depends upon
temperature.
high temp. = high pressure,
low temp. = low pressure
B. Kinetic-Molecular Theory:
1. Gases are small particles that have mass.
These particles are usually molecules, except
for the Noble Gases.
2. The particles in gases are separated
by relatively large distances.
3. The particles in
gases are in
constant rapid
motion (random).
4. Gases exert
pressure because
their particles
frequently collide
with the walls of
their container and
each other.
5. Collisions of gas particles are perfectly elastic.
Inelastic Collision
Elastic Collision
Gas particles do not slow down when hitting
each other or the walls of their container…
How do we know?
6. Temperature of a gas is
simply a measure of the
average kinetic energy
of the gas particles.
High temp. = high KE,
Low temp. = low KE
7. Gas particles exert no force on one another.
Attractive forces are so weak between particles
they are assumed to be zero.
C. Measuring Gases:
1. The following 4 variables will be used to do gas calculations:
n - amount of a gas, it is measured in moles
V - volume of a gas, it is measured in L (dm3) or mL (cm3)
“STP” = Standard Temperature & Pressure:
T- Standard Temperature: 0 oC = 273 K
Taken in oC converted to Kelvin (K)
Ex #1) 22°C =
-27.3°C =
295 K
245.7 K
T(K) = °C + 273
100.°C = 373 K
-273°C = 0 K
P - Standard Pressure: can be measured using the
following units:
(sea level pressure)
1 atm (atmospheres)
760 mm Hg
101,325 Pa (Pascals)
760 torr
101.325 kPa (kilopascals) 14.7 lb/in2 (psi)
Ex #2) Convert 1.026 atm to kPa:
 101.325 kPa 
 = 104.0 kPa
1 atm


1.026 atm 
Ex #3) Convert 98,500 Pa to mm Hg:
 760 mm Hg 
98,500 Pa 
 = 739 mm Hg
 101,325 Pa 
2. Atmospheric pressure is
the pressure exerted by
the air in the
atmosphere. This
pressure varies with
altitude and water vapor
content.
3. Atmospheric
pressure is
measured with a
barometer. This is
a glass tube sealed
at one end and
filled with Hg.
Why toxic mercury
and not water?
Atmospheric Pressure
• Affects Weather:
• Changes with Altitude:
Boiling Point Changes
w/Altitude & Pressure
12.2 The Gas Laws
A. Boyle’s Law: Pressure - Volume
Relationship. The pressure & volume of a
sample of gas at constant temperature are
inversely proportional to each other.
Indirect
P1V1 = P2V2
Doubling Pressure = ½ Volume
Quadrupling Pressure = ¼ Volume
Boyle’s Law
Ex #4) A gas has a volume of 300. mL under a pressure
of 740. mm of mercury. If the temperature remains
constant, calculate the volume when under a pressure
of 750. mm Hg.
P1V1 = P2V2
 740. mm Hg  300. mL  = 750. mm Hg   V2 
 740. mm Hg  300. mL 
750. mm Hg
=
 750. mm Hg   V2 
750. mm Hg
296 mL = V2
B. Charles’ Law: Temperature - Volume
Relationship. At constant pressure the volume
of a fixed amount of gas is directly
proportional to its absolute temperature.
Direct
V1
V2
=
T1
T2
*Temperatures must be in Kelvin!
K = °C + 273
Ballon in cool and cold water:
Ex #5)
What is the Celsius
temperature of
68.20 mL of
methane, if it
occupies a volume
of 0.02200 L at
50.0°C?
V1 V2
=
T1
T2
68.20 mL 22.00 mL
=
T1
323.0 K
 68.20 mL 323.0 K  = T1   22.00 mL 
22.00 mL
22.00 mL
T1 = 1001 K
-273 C
728 C
C. Avogadro’s Law: Amount - Volume Relationship.
Equal volumes of gases at the same temperature and
pressure contain an equal number of particles.
constant
volume
22.4L
4 He molar mass 222 Rn
1 mole gas = 22.4 L = 6.02 x 1023 particles at
STP (273 K & 1 atm)
He
O2
Rn
4 g/mol
32 g/mol
222 g/mol
Therefore because of Avogadro’s Law if these
three gases have the same number of particles
and are at the same temperature and pressure,
they must take up the same volume.
Molar Mass does not affect
volume of a gas
Dalton’s Law of Partial Pressures
D. Dalton’s Law of Partial Pressures: The sum of the
partial pressures of all the components in a gas mixture
is equal to the total pressure of the gas mixture.
PT = Pa + Pb + Pc + …
Ex #6) A flask contains a mixture of oxygen, argon, and carbon
dioxide with partial pressures of 745 torr, 0.278 atm, and
391 torr respectively. What is the total pressure in the flask?
 760 torr 
.278 atm 
 = 211 torr
 1 atm 
+ 745 torr
+ 391 torr
1347 torr
Ex #7) The total pressure of a mixture of helium
and neon is 498 mm Hg. If helium is 20.0 % of
the mixture, what is the partial pressure of
helium in mm Hg?
498 mm Hg  0.200 He   99.6 mm Hg for He
OR
20.0
x
=
100. 498 mm Hg
x = 99.6 mm Hg
E. The Combined Gas Law:
“Choyles” This law can be
used to determine how
changing two variables at a
time affects a third variable.
P1V1 P2V2
=
T1
T2
Ex #8) A gas occupies 72.0 mL at 25 °C and 198 kPa.
Convert these to standard conditions. What is the
PV
PV
=
new volume?
T
T
P1 = 198 kPa
P2 = 101.325 kPa
V1 = 72.0 mL
V2 = ?
T1 = 298 K
T2 = 273 K
1
1
1
2
2
2
(273 K) 198 kPa  72.0 mL 
(101.325 kPa)298 K
101.325 kPa  V2

=
129 mL = V2
273 K
12.3 A. Ideal Gas Law
Although no “ideal gas” exists, this law can be
used to explain the behavior of real gases
under ordinary conditions.
P = pressure (atm)
PV = nRT
V = volume (L or dm3)
n = number of moles
Deflategate 2014 – How we
R = 0.0821 L•atm/mol•K
know Tom Brady Cheated
universal gas constant
T = Kelvin temperature
B. Ideal Gas Law & The Kinetic –
Molecular Theory:
Under normal conditions (temperature and
pressure) gases behave ideally.
n ____ P ____ (more gas, ____________)
more collisions
T ____ P ____ (moves faster, ______________)
more collisions
V ____P ____
more collisions
(smaller volume, ____________)
2. Gases at low temperatures and high pressures
do NOT behave ideally. As you decrease the
volume of a gas, the volume of the particles
themselves becomes significant. The
Kinectic-Molecular Theory & Ideal Gas Law
assume that gas particles have no volume of
their own.
3. Second, the attractive forces which are very
small when the particles are moving fast,
become larger as they slow down.
Ex #9) How many grams of carbon dioxide
occupy a volume of 36.9 mL at 158 kPa
and 72 oC?
PV = nRT
 1L 
36.9 mL 
 = 0.0369 L
 1000 mL 
1st convert units to “R” units
1 atm


158 kPa 
 = 1.56 atm
 101.325 kPa 
72 C + 273 = 345 K
1.56 atm 0.0369L = n  0.0821 L  atm/mol  K 345 K 
(0.0821 L•atm/mol•K) (345 K)
(0.0821 L•atm/mol•K) (345 K)
n = 0.00203 mol CO2
 44.0 g CO2 
0.00203 mol CO2 
 = 0.0893 g CO2
 1 mol CO2 
CFC’s “chloroflurocarbons”
A chlorofluorocarbon (CFC) is an organic
compound that contains only carbon, chlorine,
hydrogen and fluorine. They are also commonly
known by the brand name Freon. Many CFCs
have been widely used as refrigerants, propellants
(in aerosol applications), and solvents. The
manufacture of such compounds has been phased
out because they contribute to ozone depletion in
the upper atmosphere. CFCs have been banned
from production in the United States since
December 31, 1995.
C. Lifting Power of Gases:
1. Uses a gas “lighter” than air (smaller molar mass.)
Ex) The Hindenburg used hydrogen
Ex) Today’s blimps use helium
2. Hot Air balloons heat air to lower its density.
3. Effusion is the movement of gas atoms or molecules through a
hole so tiny they move one particle at a time. Smaller
particles effuse faster than larger particles.
Ex)
H2
fastest
→
He
rare/expensive
→
N2
(air) plentiful/
cheap/slowest
Effusion
Effusion
The Hindenburg was a
Dirigible (a balloon that
has engines to control its
horizontal movement)
Passenger
Quarters
had 25
two-berth
cabins
Titanic - 882 feet
Hindenberg – 804 feet
May 6, 1937 Lakehurst, NJ
Memorial in
Lakehurst, NJ
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