Real gas

Definition

A real gas is a gas that does not behave ideally and deviates from the predictions of the ideal gas law due to intermolecular forces and the volume occupied by gas particles. In contrast to an ideal gas, which assumes no interactions between particles and that they occupy no volume, real gases exhibit behaviors influenced by their molecular characteristics, especially under high pressure and low temperature conditions.

5 Must Know Facts For Your Next Test

1. Real gases tend to behave more like ideal gases at high temperatures and low pressures because the influence of intermolecular forces becomes less significant.
2. As pressure increases or temperature decreases, real gases deviate from ideal behavior due to increased interactions between particles and their finite size.
3. The Van der Waals equation helps in predicting the behavior of real gases by incorporating constants that reflect the strength of intermolecular attractions and the volume occupied by molecules.
4. Real gases can exhibit phenomena such as condensation and critical points, which cannot be explained by the ideal gas law.
5. The compressibility factor (Z) is used to measure how much a real gas deviates from ideal gas behavior, with Z = 1 indicating ideal behavior.

Review Questions

• Compare and contrast real gases with ideal gases in terms of behavior under varying temperature and pressure conditions.
  • Real gases differ significantly from ideal gases at high pressures and low temperatures. While ideal gases follow the ideal gas law without deviation, real gases experience intermolecular attractions and repulsions that cause them to behave non-ideally. At high temperatures, the kinetic energy of gas molecules overcomes these forces, making real gases approximate ideal behavior. However, at low temperatures or high pressures, these interactions become pronounced, leading to deviations from the expected predictions of the ideal gas law.
• How does the Van der Waals equation improve our understanding of real gas behavior compared to the ideal gas law?
  • The Van der Waals equation enhances our understanding of real gas behavior by accounting for two critical factors: the volume occupied by gas molecules and the attractive forces between them. Unlike the ideal gas law, which assumes no volume and no interactions, the Van der Waals equation introduces correction factors 'a' and 'b' to represent these effects. This allows for more accurate predictions of gas behavior under conditions where deviations are significant, especially during phase changes and near critical points.
• Evaluate how the concept of compressibility factors can be utilized to assess deviations in real gases from ideal behavior in practical applications.
  • The compressibility factor is a crucial tool in evaluating how real gases deviate from ideal behavior in various practical applications. By measuring Z values, engineers can predict how gases will behave in industrial processes such as chemical reactions or when transported under different conditions. A compressibility factor greater than one indicates that a gas is less compressible than predicted by the ideal model, often seen under high pressure. This understanding allows for better design and optimization of equipment like reactors and pipelines, ensuring safety and efficiency in handling gases in real-world scenarios.

Related terms

Ideal Gas Law: An equation of state for an ideal gas that relates pressure, volume, temperature, and number of moles, represented as PV = nRT.

Van der Waals Equation: An adjustment to the ideal gas law that accounts for the volume of gas molecules and intermolecular forces, expressed as (P + a(n/V)^2)(V - nb) = nRT.

Compressibility Factor: A factor that quantifies how much a real gas deviates from ideal behavior, defined as Z = PV/nRT, where Z is the compressibility factor.