Why Real Gases Act Differently Than Ideal Gases

Gases in real life don’t always behave like we expect them to based on simple rules. This happens because of two main things: the forces between gas particles and how much space the particles actually take up.

What Is the Ideal Gas Law?
The ideal gas law is shown by the equation $PV = nRT$. This rule assumes that gas particles don’t interact with each other and that they don’t take up any space. However, in the real world, gases can behave differently—especially when there’s a lot of pressure or when it’s really cold.

Intermolecular Forces
One big reason gases act differently is due to the forces between the gas molecules. In the ideal gas model, these forces are not there. But for real gases, we see some important interactions:

• Van der Waals Forces: These are weak attractions that occur when gas molecules get close together, especially under high pressure.
• Dipole-Dipole Interactions: Some molecules have areas that are more positive or negative, leading to attractions between them.
• London Dispersion Forces: These forces are present in all molecules. They are stronger in larger molecules and can also cause differences from what we expect.

When gases are at high pressure, their molecules are pushed closer together. This makes the forces between them stronger, which can change how much pressure we measure compared to what the ideal gas law predicts.

Size Matters
Real gas molecules actually take up space. The ideal gas law assumes this is unimportant. But when gases are squeezed (like under high pressure), the space the molecules need becomes important. Bigger molecules take up more space than smaller ones, so we need to think about that when we measure gases.

To better describe real gases, scientist Van der Waals created a different equation:

$$[P + a \left(\frac{n}{V}\right)^2](V - nb) = nRT$$

In this equation:

• ( P ) is the pressure of the gas,
• ( V ) is the volume,
• ( n ) is how much gas we have,
• ( R ) is a constant for gases,
• ( T ) is the temperature,
• ( a ) measures the attraction between particles,
• ( b ) takes into account the space the particles occupy.

Here’s how ( a ) and ( b ) help:

• ( a ) adjusts the pressure upwards because real gases have attractions that make them act differently than expected.
• ( b ) lowers the volume to account for the space taken up by the gas particles.

When Do Gases Act Differently?

1. High Pressure: When pressure goes up, gas particles get closer together. This means their volume becomes important, and the attractions between them can actually lower the pressure compared to what we’d expect.

2. Low Temperature: If we cool the gas down, its particles move slower and start to interact more. This can lead to the gas turning into a liquid.

3. Type of Gas: Different gases behave in different ways based on their structure:

   • Nonpolar gases (like helium) usually don’t deviate much from ideal behavior.
   • Polar gases (like water vapor) have stronger attractions and show more deviations because of that.
   • Bigger molecules tend to have stronger London dispersion forces, which also lead to bigger differences from the ideal gas behavior.

Why This Matters in the Real World
Knowing how real gases behave differently is important in engineering and other fields. For example, in chemical reactions where gases are involved, understanding these behaviors helps predict how the reactions will go.

When designing high-pressure items, like storage tanks for gases, it’s important to know about these forces to keep everything safe and efficient.

In simulations, engineers often use the Van der Waals equation to ensure that they’re getting realistic results for how gases behave under different conditions.

In summary, real gases don’t always follow the simple rules we expect because of intermolecular forces, the size of the particles, and conditions like high pressure and low temperature. The Van der Waals equation gives us a better way to understand real gas behavior than the ideal gas law does. This knowledge is not just theoretical; it's also very practical for engineers and others working with gases in various situations.