# Gases

## Read in this section

### Boyles Law and pressure/volume relationships

Boyles Law states that, for any sample of gas, its pressure multiplied by the volume it occupies is a constant providing the temperature remains constant. (ie pV = constant). This is always true for perfect gases, for which pV = nRT by definition.

### The Van der Waals Equation

How can deviations from ideal behaviour be compensated for in our model of the situation? Recall that our model produced p.Vm = RT (ideal gas equation).

### Partial pressure and Daltons Law

If we consider a mixture of gases in a container, then they each exert a pressure on the walls of that container. ie: one gas provides a component of the total pressure, and another gas provides another component.

### Real Gases

Thus far we have concentrated on perfect gases. It is important to appreciate that no gas actually is perfect – they all deviate from ideal behaviour to some degree.

### The Mean Free Path

The mean free path is, as the name suggests, the average distance a molecule can go before colliding with another molecule.

### Temperature Variation of the Maxwell Distribution

Experimentally we find that the most probable speed increases as the temperature is increased, or as the moleclular mass is decreased.

### Collision Frequency (z)

We treat the molecules as hard spheres (of diameter d) – like pool balls. For two molecules to collide, their centres must come within a distance d of each other.

### Root Mean Square Speed

The root mean square speed, crms, can be related to macroscopic properties.

### The Kinetic Model of Gases

The kinetic model is based upon 3 assumptions