Non-Ideal Gases and the Van der Waals Equation

The video discusses non-ideal gases, highlighting the assumptions made in ideal gas laws and the conditions under which these assumptions fail. It introduces the Van der Waals equation as a more accurate model for non-ideal gases, explaining its correction factors for pressure and volume. The video compares calculations using the ideal gas law and the Van der Waals equation, demonstrating the importance of using the latter under certain conditions.

What is one of the assumptions made about ideal gases in kinetic molecular theory?

• Gas particles move at very low speeds.

• Gas particles are dimensionless points.

• Gas particles interact strongly with each other.

• Gas particles have significant volume.

Under what conditions do gases deviate from ideal behavior?

• High temperatures and high pressures

• Low temperatures and low pressures

• Low temperatures and high pressures

• High temperatures and low pressures

What does the Van der Waals equation account for that the ideal gas law does not?

• The speed of gas particles

• The shape of gas particles

• The volume and interactions of gas particles

• The color of gas particles

How does the Van der Waals equation modify the pressure term?

• By adding a correction factor related to particle attraction

• By adding a constant to the pressure

• By subtracting a constant from the pressure

• By multiplying the pressure by a constant

In the example given, what was the pressure calculated using the ideal gas law?

• 31.2 atmospheres

• 34.6 atmospheres

• 33.5 atmospheres

• 32.4 atmospheres

Why is the Van der Waals equation more accurate for carbon dioxide in the example?

• Because carbon dioxide is a solid at room temperature

• Because carbon dioxide has a high boiling point

• Because carbon dioxide is a nonpolar molecule

• Because carbon dioxide is a polar molecule

What would make the deviation between the ideal gas law and Van der Waals equation more significant?

• Increasing the pressure to hundreds of atmospheres

• Lowering the temperature further

• Using a smaller container

• Decreasing the number of moles