Boyle's Law fails at low temperatures and high pressures because real gases deviate from the ideal gas assumptions that the law relies on.
Read also Derivation of Boyle's Law from Kinetic Gas Equation
Key Reasons for Failure
1. Intermolecular Forces Become Significant
- Boyle's Law assumes no attractive or repulsive forces between gas molecules.
- At high pressures, molecules are forced closer together, making intermolecular attractions significant.
- At low temperatures, molecules slow down, allowing weak intermolecular forces to have greater effect on their movement.
- These forces pull molecules together, so pressure increases less than predicted by Boyle's Law.
2. Molecular Volume Becomes Important
- Boyle's Law assumes molecules are point particles with no volume.
- In reality, molecules have actual size that occupies measurable space.
- At high pressures when compressed, the space molecules take up becomes a significant part of the total volume.
- This makes the pressure-volume relationship no longer perfectly inverse.
3. Gas Liquefaction
- At low temperatures, gas molecules slow down and tend to stick together.
- If cooled enough, gases turn into liquids (liquefy).
- Boyle's Law assumes the gas stays in the gaseous state and doesn't account for condensation.
- Once the gas condenses to liquid, the law no longer applies.
Summary Table
| Condition | What Happens | Why Boyle's Law Fails |
|---|---|---|
| High Pressure | Molecules forced close together | Intermolecular forces + molecular volume become significant |
| Low Temperature | Molecules slow down, stick together | Intermolecular forces dominate; gas may liquefy |
Key Takeaway: Real gases behave ideally only at low pressures and moderate to high temperatures.