Behaviour of Real Gases — Definition

Definition

Imagine you have a bunch of tiny, invisible particles zipping around in a container. If these particles were perfectly tiny, like points, and never bumped into each other except for perfectly elastic collisions with the container walls, and never felt any attraction or repulsion towards each other, then you'd have what we call an 'ideal gas'.

This is a theoretical concept, a perfect model that helps us understand gas behavior under very specific, simplified conditions. The ideal gas equation, $PV = nRT$, works perfectly for this imaginary gas.

Now, let's talk about 'real gases' – these are the actual gases we encounter every day, like oxygen, nitrogen, carbon dioxide, or methane. These real gases are not perfect. Their molecules are not infinitely small; they actually take up some space, even if it's tiny.

Also, these molecules aren't completely indifferent to each other; they do experience weak attractive forces (like van der Waals forces) and sometimes repulsive forces when they get too close. Because of these two fundamental differences – finite molecular volume and intermolecular forces – real gases don't always behave exactly like ideal gases.

Think of it this way: at high temperatures and low pressures, real gas molecules are far apart and moving very fast. In this scenario, their own volume is negligible compared to the total volume of the container, and the attractive forces between them are too weak to have a significant effect because they rarely get close enough. So, under these conditions, real gases behave *almost* ideally, and the $PV = nRT$ equation works quite well.

However, when you increase the pressure, the gas molecules get pushed closer together. Now, their actual volume starts to become a noticeable fraction of the total volume, and the attractive forces between them become much stronger because they are in closer proximity.

Similarly, if you lower the temperature, the molecules move slower, allowing the attractive forces to 'catch' them more effectively, pulling them closer and reducing the pressure they exert on the container walls.

In these conditions (high pressure, low temperature), real gases deviate significantly from ideal behavior. The pressure they exert might be less than predicted by the ideal gas law (due to attractions), or the volume they occupy might be more than predicted (due to their own finite volume).

The van der Waals equation is a famous attempt to correct the ideal gas law for these real gas imperfections.