properties of gases

Gases : Kinetic Molecular Theory

The Kinetic Molecular Theory is used to explain the behavior of gases and is based upon the following postulates:

Gases are composed of a many particles that behave like hard spherical objects in a state of constant, random motion.

These particles move in a straight line until they collide with another particle or the walls of the container.

These particles are much smaller than the distance between particles, therefore the volume of a gas is mostly empty space and the volume of the gas molecule themselves is negligible.

There is no force of attraction between gas particles or between the particles and the walls of the container.

Collisions between gas particles or collisions with the walls of the container are elastic. That is, none of the energy of the gas particle is lost in a collision.

The average kinetic energy of a collection of gas particles is dependent only upon the temperature of the gas.

In short, the key parts are:

constant random motion
straight line motion
negligible volume
no forces of attraction
elastic collisions

Check out this Animated Molecular Model for an Ideal Gas!

Kinetic Energy - The energy of motion.

where m is mass and v is velocity
Systems as a whole are described by an average kinetic energy.

Using the Kinetic Molecular Theory to explain the Gas Laws

The Relationship Between P and n
Boyle's Law
Amonton's Law
Charles' Law
Avogadro's Hypothesis
Dalton's Law of Partial Pressures

The Relationship Between P and n

P is the force exerted on the walls of the container during a collision
An increase in the number of particles increases the frequency of collisions with the walls
Therefore P increases as n increases.

Boyle's Law

Compressing a gas makes the V smaller but does not alter the KE_avgof the molecules since T is constant.
Though the speed of the particles remains constant, the frequency of collisions increases because the container is smaller.
Therefore, P increases as V decreases.

Amonton's Law

The KE_avg of a gas particle increases as T increases.
Since the mass remains constant, the average velocity of the particles must increase (KE = 1/2mv²).
As velocity increases, so does force, thus the force per collision increases as T increases.
P = force/area, therefore P increases as T increases.

Charles' Law

The KE_avg of a gas particle is proportional to T.
Since mass is constant, the average velocity of the particles must increase (KE = 1/2mv²).
At higher velocity, the particles exert greater force which increases P.
If the walls are flexible, they will expand to balance the atmospheric pressure outside.
Therefore, V is directly proportional to T.

Avogadro's Hypothesis

An increase in the number of particles increases the frequency of collisions with the walls, thus P increases.
In a flexible container, the walls will expand until the pressure of the gas equals the atmospheric pressure outside the container.

Dalton's Law of Partial Pressures

If ball bearings of different sizes were placed in a moving container, the total number of collisions between the balls and the walls would be equal to the sum of the collisions that would occur if each ball were present by itself in the container.
Therefore, the pressure of a mixture of gases is equal to the sum of the partial pressures of the individual gases.

Graham's Laws of Diffusion and Effusion