Question

thermodynamics:

why can’t the van der Waaks equation extend itself to the low temperature regions near absolute zero ?

Answer 1

At high pressures and low temperatures, intermolecular forces between gas particles can cause significant deviation from ideal behavior.

Key Points

• Ideal gases are modeled as interacting through perfectly elastic collisions, implying that intermolecular interactions do not significantly contribute to the gas particles’ energetics.
• Real gas interactions, such as attractive and repulsive intermolecular forces, are more complex than perfectly elastic collisions; the significance of these contributions varies with the gases’ conditions.
• The van der Waals equation takes into account these intermolecular forces and offers an improved model for real gas behavior.

The Ideal Gas Law is a convenient approximation for predicting the behavior of gases at low pressures and high temperatures. This equation assumes that gas molecules interact with their neighbors solely through perfectly elastic collisions, and that particles exert no intermolecular forces upon each other.

Intermolecular forces describe the attraction and repulsion between particles. They include:

• Dipole -dipole forces
• Ion-dipole forces
• Dipole-induced dipole forces or Debye forces
• Instantaneous dipole-induced dipole forces or London dispersion forces.

The contribution of intermolecular forces creates deviations from ideal behavior at high pressures and low temperatures, and when the gas particles’ weight becomes significant.

• At low temperatures, gas particles have less kinetic energy, and therefore move more slowly; at slower speeds, they are much more likely to interact (attracting or repelling one another) upon collision. The Ideal Gas Law does not account for these interactions.

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