GAS LAWS
The Nature of Gases
Gases expand to fill their containers
Gases are fluid – they flow
Gases have low density
1/1000 the density of the equivalent liquid or solid
Gases are compressible
Gases effuse and diffuse
Kinetic Theory
• Gases are composed of small, separate particles called
molecules
• Move in constant motion
• All collisions between particles are perfectly elastic
• The molecules of a gas display no attraction or repulsion
• The average kinetic energy of the molecules is directly proportional
to Kelvin temperature of the gas
Ideal Gases
Ideal gases are imaginary gases that perfectly
fit all of the assumptions of the kinetic molecular
theory.
 Gases consist of tiny particles that are far apart
relative to their size.
 Collisions between gas particles and between
particles and the walls of the container are
elastic collisions
 No kinetic energy is lost in elastic collisions
Ideal Gases (continued)
 Gas particles are in constant, rapid motion. They
therefore possess kinetic energy, the energy of
motion
 There are no forces of attraction between gas
particles
 The average kinetic energy of gas particles
depends on temperature, not on the identity
of the particle.
Gas Pressure
• Pressure= Force/Area
• Atmospheric pressure:
• The pressure the earth’s atmosphere exerts due to its weight
• Barometer:
• Instrument used to measure atmospheric pressure
• Invented by Toricelli
• Baro= weight
• Meter= measure
• Normal Atmospheric Pressure:
• Also called standard pressure
• 1 atm = 760 mm Hg = 760 torr
• 1 atm= 101.3 kPa
Gas Pressure
• Barometer:
• Instrument used to measure
atmospheric pressure
• Invented by Toricelli
• 760 mm Hg at 1 atm
• Normal Atmospheric
Pressure:
• Also called standard pressure
• 760 mmHg
• 760 torr
• 1 atm
• 101.3 kPa
• 273 K
Gas Pressure
• STP
• Standard temp and pressure
• P= 1 atm, 760 torr
• T= 0ºC, 273 K
• Molar volume of ideal gas is 22.4 L at STP
• Manometer
• Instrument used to measure gas pressure
• U-shaped tube partially filled with mercury
• One end open to confined gas
• One end open to atmosphere
Converting Celsius to Kelvin
Gas law problems involving temperature require
that the temperature be in KELVINS!
Kelvins = C + 273
°C = Kelvins - 273
Boyle’s Law
• When temperature is held constant, the pressure and
volume of a gas are inversely proportional
• If P goes up, V goes down
Charles’s Law
• When pressure is held constant, the volume and
temperature of a gas are directly proportional
• If V goes up, T goes up
Gay Lussac’s Law
• When volume is held constant, the pressure and
temperature of a gas are directly proportional
• If P goes up, T goes up
The Combined Gas Law
The combined gas law expresses the
relationship between pressure, volume and
temperature of a fixed amount of gas.
P1V1 P2V2

T1
T2
Boyle’s law, Gay-Lussac’s law, and Charles’ law
are all derived from this by holding a variable
constant.
Health Note
When a scuba diver is several
hundred feet under water, the
high pressures cause N2 from the
tank air to dissolve in the blood.
If the diver rises too fast, the
dissolved N2 will form bubbles in
the blood, a dangerous and
painful condition called "the
bends". Helium, which is inert,
less dense, and does not dissolve
in the blood, is mixed with O2 in
scuba tanks used for deep
descents.
Dalton’s Law of Partial Pressures
• Each gas exerts the same pressure it would if it alone was
present at the same temperature
• Gas collected over water- pressure in the container is the
sum of the vapor pressure of the gas and the water’s
vapor pressure
• Subtract the water vapor pressure from the total pressure
to obtain the pressure of the gas alone
Solve This!
A student collects
some hydrogen
gas over water at
20 degrees C
and 768 torr.
What is the
pressure of the
H2 gas?
768 torr – 17.5 torr = 750.5 torr
Ideal Gas Law
PV = nRT
P = pressure in atm
V = volume in liters
n = moles
R = proportionality constant
= 0.08206 L atm/ mol·K
T = temperature in Kelvins
Holds closely at P < 1 atm
Avogadro’s Law
• Equal volumes of different gases, at the same temp and
pressure contain the same number of molecules
• How would the number of molecules in 2 liters of
hydrogen gas compare with the number of molecules in 2
liters of oxygen gas at the same temperature and
pressure?
• Why is 22.4 liters called the molar volume of gas?
Avogadro’s Law
Equal volumes of gases at the same T
and P have the same number of
molecules.
V = n (RT/P)
V and n are directly related.
twice as many
molecules
Gas Density
mass
molar mass
Density 

volume molar volume
… so at STP…
molar mass
Density 
22.4 L
Density and the Ideal Gas Law
Combining the formula for density with the Ideal
Gas law, substituting and rearranging algebraically:
MP
D
RT
M = Molar Mass
P = Pressure
R = Gas Constant
T = Temperature in Kelvins