The kinetic theory of gases is an attempt to explain the behavior
of gases in terms of atoms and molecules. It relies on a set of postulates
that seem wordy, but actually include a remarkable amount of detail about
gases.
- Gases are made of atoms and molecules in
constant, random motion. These particles collide with each other and
with the walls of the container the gas is held in.
- Collisions between the gas particles and
the walls causes pressure.
- Collisions between gas particles are elastic:
no energy is gained or lost in such a collision.
- The volume of the particles in a gas is
negligible compared to the volume of the container.
- There are no attractive forces between
atoms or molecules in the gas
- The average translational kinetic energy
of the particles in a gas is proportional to the temperature of the
gas. (Translational means movement: many gas molecules are also
rotating or vibrating, but this doesn't affect the temperature in this
model.)
- At a given temperature, all gases have
the same average translational kinetic energy.
Some of these are clearly only approximations:
for example, we all know that atoms and molecules actually do have a
volume. However, at room temperature and pressure, this is a reasonable
approximation. (The kinetic theory can be patched up to more closely model
real gases.) |