Real gases deviate from ideal behavior mainly because the assumptions of the ideal gas law don’t perfectly apply to them. Here’s why and how:
- Intermolecular forces exist
- Ideal gases assume no attraction or repulsion between particles.
- In real gases, particles attract each other slightly, especially at low temperatures. This makes the gas less “spread out” than predicted, so it behaves differently from an ideal gas.
- Finite particle volume
- Ideal gases assume gas particles have negligible size.
- Real gas molecules take up space, so at high pressures (when particles are close together), the volume of the gas is larger than predicted by the ideal gas law.
- Low temperature and high pressure effects
- At low temperatures, particles move slower, so attractions become more significant, and the gas can even condense into a liquid.
- At high pressures, particles are forced close together, making their volume and interactions important, causing deviations from ideal behavior.
In short: Real gases deviate from ideal behavior mainly at high pressure and low temperature, due to intermolecular forces and finite molecular size.