Real Gas Behaviour & Ideality Limitations
This is Higher Level (HL) content.
No real gas behaves ideally under all conditions. Deviations occur overwhelmingly under extreme conditions: specifically, low temperature and high pressure.
Why Real Gases Deviate
Low Temperature: Intermolecular Forces Engage
At low temperatures, particle kinetic energy decreases significantly. Particles move sluggishly and intermolecular attractive forces are no longer negligible. Particles pull on one another, colliding with container walls with reduced force: observed pressure is lower than ideal predictions (\(p_{\text{real}} < p_{\text{ideal}}\)).
High Pressure: Molecular Volume Non-Negligible
Under immense pressure, particles are forced into close proximity. The volume occupied by the particles themselves becomes a significant fraction of the total space. The actual volume available for motion is reduced: measured volume is greater than ideal predictions (\(V_{\text{real}} > V_{\text{ideal}}\)).
When Does a Gas Behave Most Ideally?
A gas behaves most like an ideal gas at high temperature (particles have high kinetic energy, overcoming intermolecular attractions) and low pressure (particles are widely separated, rendering particle volume negligible).
Compressibility Factor (Z) vs Pressure
Summary: Which Assumption Breaks Down?
| Condition | Assumption That Breaks Down | Observed Effect |
|---|---|---|
| Low temperature | "No intermolecular forces" | Preal < Pideal |
| High pressure | "Negligible particle volume" | Vreal > Videal |
Which Gases Deviate the Most?
Molecular Mass & Intermolecular Force Effects
- Gases closest to ideal: Small, non-polar molecules with weak London dispersion forces (e.g. Helium, Hydrogen).
- Gases deviating most: Large molecules with strong London dispersion forces, and polar molecules with dipole-dipole or hydrogen bonding (e.g. NH₃, H₂O, CO₂).
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