Short Answer:
Real gases deviate from ideal gases mainly at high pressures and low temperatures, where the assumptions of the ideal gas law no longer hold true. In real gases, molecules have finite volume and experience intermolecular forces, which are not considered in the ideal gas model.
Due to these factors, the behavior of real gases becomes different from that predicted by the ideal gas law. Real gases can show higher or lower pressure than expected. To correct these deviations, engineers use the compressibility factor (Z) or other equations like the van der Waals equation.
Detailed Explanation:
Real gases deviate from ideal gases
The ideal gas law (PV = nRT) is based on a simplified model where gas particles are treated as point masses with no interactions. This model works well for many gases under normal conditions (room temperature and atmospheric pressure). However, in practical applications, gases often operate under high pressure, low temperature, or in dense conditions, where they show behavior that is different from ideal gases.
These differences are known as deviations, and understanding them is important for accurate gas calculations in engineering systems like compressors, turbines, pipelines, and refrigeration.
Reasons for Deviation of Real Gases
- Molecular Volume
- Ideal gases assume gas molecules have no size.
- But real gas molecules have finite volume.
- At high pressures, molecules are very close, and their own volume becomes significant, reducing the free space for movement.
- Intermolecular Forces
- Ideal gases assume no attraction or repulsion between molecules.
- Real gases have van der Waals forces:
- Attractive forces pull molecules together at low pressure or low temperature, reducing pressure.
- Repulsive forces push molecules apart at very high pressure, increasing pressure.
- Non-Ideal Collisions
- In ideal gases, collisions are perfectly elastic.
- In real gases, energy may be lost or absorbed during collisions, affecting pressure and temperature.
Conditions When Deviations Are Significant
- High Pressure
- Molecules are compressed closer, increasing the effect of molecular size and repulsive forces.
- Pressure becomes higher than predicted by ideal gas law.
- Low Temperature
- Molecules move slower, and attractive forces dominate.
- Gas is more compressible, and pressure becomes lower than predicted.
- Near Condensation Point
- Gases begin to liquefy, and ideal assumptions completely fail.
Correction Using Compressibility Factor
To correct for deviation, engineers use the compressibility factor (Z):
Z = (PV) / (nRT)
- If Z = 1 → gas behaves ideally
- If Z < 1 → attractive forces dominate (gas is more compressible)
- If Z > 1 → repulsive forces dominate (gas is less compressible)
Using this factor, we modify the ideal gas law:
PV = ZnRT
This gives more accurate results for real gas behavior.
Van der Waals Equation
Another method to account for deviation is the van der Waals equation:
[P + a(n/V)²] × [V – nb] = nRT
Where:
- a = corrects for intermolecular attraction
- b = corrects for molecular volume
This equation adjusts both pressure and volume to match real gas behavior.
Practical Applications
- Refrigeration systems – where gases operate at low temperatures.
- Natural gas pipelines – where pressure is very high.
- Combustion engines – gas properties affect performance.
- Gas storage and transport – volume and pressure must be precisely calculated.
- Aerospace and chemical plants – where exact gas behavior is critical.
Conclusion
Real gases deviate from ideal gases due to molecular size and intermolecular forces, especially under extreme conditions like high pressure and low temperature. While the ideal gas law simplifies calculations, it does not always give accurate results for real gases. Engineers use compressibility factors or advanced models like the van der Waals equation to correct these deviations. Understanding how and why gases deviate from ideal behavior is essential in designing safe and efficient mechanical and thermal systems.