Structure 1.5.2—Real gases deviate from the ideal gas model

Why Real Gases Deviate from Ideal Behavior 

Real gases deviate from ideal gas behavior because the assumptions of the kinetic molecular theory are not fully valid under all conditions.

The most significant deviations occur at low temperatures and high pressures.

1. Intermolecular Forces Become Significant (Low Temperature)

• At low temperatures, gas particles move more slowly and are closer together.
• Attractive forces between molecules (e.g. Van der Waals or dipole-dipole) become more significant.
• These forces cause particles to exert less pressure on container walls than predicted by ideal gas law.
• Effect: The observed pressure is lower than the ideal gas model predicts.

2. Particle Volume Becomes Significant (High Pressure)

• At high pressure, gas particles are crowded closely together.
• Their own volume becomes significant relative to the container’s volume.
• This causes the gas to occupy more volume than ideal gas theory would predict.
• Effect: The measured volume is higher than expected.

 Compressibility Factor (Z): Quantifying Deviations

The compressibility factor \( Z \) is used to compare the behavior of a real gas to an ideal gas:

\( Z = \frac{P V_m}{R T} \)

where:

• \( P \) is pressure
• \( V_m \) is molar volume
• \( R \) is the ideal gas constant
• \( T \) is temperature in Kelvin

• If \( Z = 1 \) → the gas behaves ideally.
• If \( Z < 1 \) → gas is more compressible than ideal (intermolecular attractions dominate).
• If \( Z > 1 \) → gas is less compressible than ideal (repulsions or finite particle volume dominates).

Real-World Insight: Scientists use more accurate equations, such as the Van der Waals equation, to describe real gas behavior. These take into account particle volume and intermolecular forces.