Real gases and van der Waals equation
Real gases differ from ideal gases at high pressure and low temperature because their molecules take up space and attract each other. The van der Waals equation adds two corrections: a for attraction and b for molecule size.
Last updated
Analogy
A crowded train: people (molecules) have real size, and at rush hour they cannot move as freely as the ideal model says.
Example
(P + an²/V²)(V − nb) = nRTa corrects for attractionb corrects for molecular volume
In practice
Carbon dioxide deviates strongly from ideal behaviour near room temperature at high pressure, which is why it can be liquefied in fire extinguishers.
Practice
Flashcards
- When are real gases least ideal?At high pressure and low temperature.
Questions people ask
When does a real gas behave most like an ideal gas?
A real gas is most ideal at low pressure and high temperature. At low pressure, molecules are far apart, so their size and attraction hardly matter. At high temperature, they move so fast that attractions have little effect. Small, nonpolar gases like helium behave ideally over the widest range.
Why is the pressure of a real gas lower than ideal at moderate pressures?
At moderate pressure, attractions between molecules slow their approach to the wall and soften their collisions. This makes the measured pressure smaller than the ideal gas law predicts. At very high pressure, molecular volume takes over and the real pressure rises above the ideal value.