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### Course: Chemistry library>Unit 8

Lesson 2: Kinetic molecular theory

# The kinetic molecular theory of gases

The kinetic molecular theory (KMT) describes the behavior of ideal gases at the particle level. The five main postulates of the KMT are as follows: (1) the particles in a gas are in constant, random motion, (2) the combined volume of the particles is negligible, (3) the particles exert no forces on one another, (4) any collisions between the particles are completely elastic, and (5) the average kinetic energy of the particles is proportional to the temperature in kelvins. Created by Sal Khan.

## Want to join the conversation?

• I don't really get about the velocity of the particles being preserved when they bounce of the walls.
Could someone explain?
• When particles collide with each other, or with the sides of the container, they realistically transfer some of their kinetic energy. Kinetic energy again being the energy associated with motion and is directly proportional to a particle’s velocity. This transferred kinetic energy is transformed into other forms of energy like potential energy or heat. However in the kinetic theory of gases we assume that these collisions are what is known as elastic where the collisions do not result in the transfer of any kinetic energy and thus the particles would maintain their velocity. Now like I said realistically there is at least some kinetic energy transferred when gas particles collide so rarely are these collisions ever perfectly elastic. However when averaged over the many trillions of gas particles, these collisions can be assumed to be essentially elastic which makes the math for calculations simpler.

All these assumptions describe gases as ideal gases which obey simple equations like the ideal gas law. And while no gas is ever completely ideal, these assumptions allow us to make roughly accurate predictions on the behaviors of gases.

Hope that helps.
• How do the particles exert no force on one another? Newton's law of universal gravitational force proves the opposite. I mean, they must apply force on each other, even if it is tiny or negligible.
• He's referring to ideal gas. And that's one reason why no gas is ideal.
• What are particles made of?
(1 vote)
• Which particles are you referring to? Many things can be referred to as particles, so you'll need to be more specific.
• Can't we measure gasses in 'microscopic level' or what if we try to do that?
• No, because gas is not just one atom or molecule it is multiple. For example Oxygen can't stand alone it is O2 which is more than one atom or molecule. By studying only one atom or particle we can't measure the behavior of the gas as a whole.
• So, according to Axiom 3, the London Dispersion Forces shouldn't be there. But that's not true. So, is this right ? If not, why so ?
(1 vote)
• These axioms essentially assume that the gas behaves ideally which involves the gas particles having no attractions to each other. So using the ideal gas law: PV = nRT, you are doing so under this simplification. Of course we know things like London dispersion forces exist which does cause attraction which makes the ideal gas law only an approximation. This approximation holds though for most real world situations so the assumption is acceptable. Other gas laws like the Van Der Waals equation include corrections to account for these attractive forces.

Hope that helps.