Physical Properties of Gases
Chapter 21
Particle model of gas
behaviour:
The low density of gas, relative to a liquid
and solid, suggests that the particles of gas
are much more widely spaced. This is
consistent with the observation that gases
are easily compressed.
The observations that gases spread to fill
the space available suggests that the
particles of a gas move independently of
each other.
The wide spacing of particles together with
their movement explains why gases mix rapidly.
Kinetic molecular theory
The model used by scientists to explain gas behaviour is known as the kinetic molecular theory of
gases.
According to this model:
 gases are composed of small particles (atoms or molecules). The total volume of the
particles in the sample is very much smaller than the volume occupied by the gas. Most of
the volume taken up by a gas is empty space
 these particles move rapidly in a random, straight line motion. Particles will collide with each
other and with the walls of the container
 the bonding forces between particles are extremely weak. It is assumed that particles move
around independently
 collisions between particles are elastic, i.e. energy is conserved. Kinetic energy (energy of
movement) can be transferred from one particle to another, but the total kinetic energy will
remain constant
 the average kinetic energy of the particles increases as the temperature of the gas is increased.
Relationship between molecular kinetic energy and temperature
According to the kinetic molecular theory, the average kinetic energy of gas particles is proportional to the temperature of the
gas sample. It is important to understand, however, that, at any given temperature:
 the average kinetic energy of the gas is not dependent on the particular gas
 within each gas sample, there will be some particles of very low kinetic energy and some of high energy. The
temperature merely reflects the average kinetic energy of particles in the sample.
The distribution of the kinetic energies of particles in a gas at a given
temperature can be illustrated by the graph. The graph refers to the
same sample of oxygen gas. However, it could equally apply to any
other gas at the stated temperatures. It shows the relative number of
oxygen molecules with a given kinetic energy at three different
temperatures.
We can see that:
 only a small proportion of molecules has very low or very
high kinetic energy

at all three temperatures, there are some molecules with very
low kinetic energy
 the proportion of molecules with high kinetic energy increases with temperature
 the average kinetic energy of the sample increases with temperature
 the area under each graph represents the total number of oxygen molecules. As the amount of gas is the same at all three
temperatures, the area under each graph must be the same.
The average kinetic energy of particles in gases is related to their average speed of movement by the relationship:
where m is the mass of the gas particles and v is the average velocity of the particles.
As the average molecular kinetic energy of particles in different gas samples is the same at a given temperature, we can deduce
that the lighter the gas atoms or molecules in a sample, the greater their average velocity.
Diffusion
Diffusion is the term used to describe the way each gas in a mixture of gases spreads itself evenly to fill the total volume
available. The rate at which gases diffuse depends on the average velocity of their particles. Therefore, gases of lower relative
molecular mass will diffuse more rapidly than gases of higher relative molecular mass at the same temperature. Diffusion occurs
more rapidly at higher temperatures than at lower temperatures. Because the average kinetic energy of the particles is greater at
higher temperatures, so is their average velocity.
Using the kinetic molecular theory
This model of gases helps to explain why strong smells spread quickly. The particles that make up the vapour or gas are in
constant motion and continue to move in all directions. Because there is so much space between these particles and the other gas
particles in the air, all the gas particles rapidly mix with each other.
The model also explains why gases expand to fill any container. The gas particles continue to move outwards until stopped by
the walls of the container. This means that the amount of space a fixed mass of gas takes up (its volume) can be altered by
changing the size of its container. This makes gases very different from liquids because a fixed mass of liquid can change its
shape but not its volume
A gas can be compressed by reducing the volume of its container because there is so much space between the gas particles. This
is a very useful property of gases and explains why compressed air can be used to inflate tyres. The more a gas is compressed,
the greater the number of collisions the gas particles will have with each other and the walls of its container. These collisions
produce a force on the walls of the container, such as the inside of a tyre, which can be measured. This force on the wall area is
measured as pressure.
Pressure
In terms of the kinetic theory of gases, pressure is defined as the force exerted on a unit area of a surface by the particles of a gas
as they collide with each other and the walls of a container. The gas pressure exerted depends on the number of collisions
between the molecules and the walls of the container.
The pressure of a fixed amount of gas is independent of the actual gas being
considered. In a gaseous mixture, such as air, the nitrogen molecules collide with
the walls of a container, exerting a pressure. In a similar way, the oxygen
molecules exert a pressure as do molecules of carbon dioxide, argon and so on
for each gas present in the mixture. The measured air pressure is the sum of these
individual gas pressures.
Partial pressure
The pressure exerted by the individual gases in a mixture is known as a partial
pressure.
In a sample of air in a container, the partial pressure of the nitrogen is the
pressure that amount of nitrogen would exert if it were the only gas in the
container.
Similarly, the partial pressure of oxygen is the pressure that amount of oxygen
would exert if it were the only gas in the container.
When nitrogen and oxygen are in the same container, the molecules of each
move independently and collide with the walls as though the other gas were not
there. The total effect of all collisions in this mixture is the sum of the effects of
the collisions of the individual gases.
The total pressure of a gas mixture is therefore the sum of the individual partial
pressures of the gases in the mixture.
The pressure will increase if the amount of gas is increased, the temperature of
the gas is increased or the volume of the container is decreased.
Key Questions:
Q1, 2, 3 pg 360