S1.5.2 Deviations from ideal gas behavior Notes

Limitations of the Ideal Gas Model and Deviations of Real Gases

1. You're inflating a balloon on a chilly winter day.
2. You notice that the balloon doesn’t expand as much as it would on a warm day, even though you’ve filled it with the same amount of air. Why is this?

The answer lies in the behavior of gases under different conditions.

Why Do Real Gases Deviate from Ideal Gas Behavior?

1. The ideal gas model is based on several simplifying assumptions, as explored in the previous section:
   1. Gas particles have negligible volume compared to the volume of their container.
   2. There are no intermolecular forces between gas particles.
   3. Collisions between gas particles are perfectly elastic.
   4. The kinetic energy of gas particles is directly proportional to their temperature in kelvin.
2. While these assumptions work well for many gases under standard conditions, they break down under specific circumstances.
3. Real gases deviate from ideal behavior primarily due to:
   1. Intermolecular Forces: At low temperatures, attractive forces between particles become significant.
   2. Finite Particle Volume: At high pressures, the volume occupied by gas particles themselves is no longer negligible.

1. Low Temperatures: The Role of Intermolecular Forces

1. At low temperatures, gas particles move more slowly because their kinetic energy decreases.
2. This reduced motion allows intermolecular forces, such as van der Waals forces, to become more pronounced.
3. These forces cause gas particles to attract each other, reducing the pressure exerted by the gas compared to what the ideal gas law predicts.