Diffusion is the term used to describe the mixing of gases. Diffusion is described in the lab, Diffusion.
In the 19th century, Thomas Graham found that the rate of effusion of a gas is inversely proportional to the square root of the mass of its particles. Graham's law of effusion is as follows: Rate of effusion for gas 1 / Rate of effusion for gas 2 = sq rt. M2 / sq rt. M1, where M1 and M2 represent the atomic or molecular weights of the gases as appropiate.
So far in this chapter, we have been talking about ideal gases, that behave according to the ideal gas law. In reality, gases do not believe this way. The postulates of the kinetic molecular theory are not all correct, as gases actually do have volume, and they are also attracted to each other. It is important to know that real gases tend to behave most like ideal ones when they are at low pressure and high temperature. In 1873, a physics professor named Johannes van der Waals developed an equation for real gases. Based on the ideal gas law, van der Waals added in correction factors. His equation, known as the van der Waals equation, is very complex and is represented:
[Pobs + a(n/V)2](V-nb) = nRT
The value a(n/V)2 is used to correct for forces of attraction. The value nb is used to correct for volume. The values for a and b are experimentally calculated for each different gas. Some of them are represented in the table below. The work of van der Waals won him the Nobel Prize in 1910.
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Unit 2 - Section 7