Using the kinetic molecular theory, explain what the difference is between a gas that exerts a pressure of 1.0 atm and one that exerts a pressure of 1.5 atm. *

In the kinetic molecular theory, the molecules of an ideal gas are in constant random motion inside the container of the gas, and the pressure of the gas (which is the pressure exerted by the molecules in their collisions with the walls of the container) arise from this random motion of the molecules.

The main assumptions of the kinetic theory of gases are:

The gas consists of a large number of molecules that collide between each other and the walls of the container; all these collisions are elastic
The duration of the collisions is negligible compared to the time between the collisions
The number of molecules is so large that statistics can be applied
Intermolecular forces between the molecules are negligible (except during the collisions)
The volume of the molecules is negligible compared to the volume of the container

In particular, the pressure of the gas is directly proportional to the average kinetic energy of the molecules, according to the equation:

[tex]pV=\frac{2}{3}K[/tex]

where

p is the pressure of the gas

V is the volume of the container

K is the average kinetic energy of the molecules in the gas

We see that as the pressure is higher, the higher the kinetic energy of the particles: this means that the molecules will move faster, on average.

Therefore in this problem, the gas that exerts a pressure of 1.5 atm will have molecules moving faster than the molecules of the gas exerting a pressure of only 1.0 atm.