Real gases can behave differently from the ideal gases we often learn about. This can greatly affect how we use them in real-life situations.

Here are some key points to understand:

• Molecular Forces: In real gases, the molecules pull towards each other or push away from each other. This can change their behavior. For example, when gases are at high pressure, their molecules get pushed closer together. This stronger interaction changes how much space the gas takes up, which the ideal gas law doesn’t consider.

• Space Taken Up by Molecules: The ideal gas law assumes that gas molecules don’t have any size. But in reality, especially when gases are at low temperatures, the size of these molecules matters. This can make the volume of the gas larger than what we’d expect if we only used the ideal gas law.

• Conditions That Aren’t Ideal: When gases are at low temperatures and high pressures, the ideal gas law does not work well. For these conditions, we use a different equation called the Van der Waals equation:

$\left(P + \frac{a}{V_m^2}\right)(V_m - b) = RT$

In this equation, $a$ represents the forces between molecules, and $b$ shows the size of the molecules.

Understanding how real gases work is important for things like burning fuel, refrigeration, and moving gas. If engineers know about these differences, they can design better systems that work more efficiently.

When doing calculations, using the Van der Waals equation instead of the ideal gas law can lead to more accurate results. This is especially important in engineering, where knowing how gases behave correctly can affect how well a system works.