Why Real Gases Deviate from Ideal Behaviour
Real gases are known to deviate from ideal gas behavior under certain conditions. This deviation is a result of several factors that affect the behavior of real gases. Understanding why real gases deviate from ideal behavior is crucial in various scientific and engineering applications, such as the design of gas compressors, refrigeration systems, and chemical processes. In this article, we will explore the reasons behind these deviations and their implications.
One of the primary reasons why real gases deviate from ideal behavior is the intermolecular forces between gas molecules. Ideal gases are assumed to have no intermolecular forces, which means that they do not interact with each other. However, real gases have attractive and repulsive forces between their molecules. These forces become more significant at higher pressures and lower temperatures, leading to deviations from ideal gas behavior.
Another factor that contributes to the deviation is the finite volume of gas molecules. In the ideal gas model, gas molecules are considered to be point particles with no volume. In reality, gas molecules occupy a certain volume, which becomes more significant at higher pressures. This finite volume leads to a deviation from the ideal gas law, as the volume of the gas molecules affects the overall volume of the gas.
The shape of gas molecules also plays a role in the deviation from ideal behavior. Ideal gases are assumed to be spherical, but real gas molecules can have more complex shapes. This difference in shape can affect the interactions between molecules and, consequently, the pressure and volume of the gas.
The presence of a non-zero value of the Van der Waals constant ‘a’ is another reason for the deviation from ideal behavior. The Van der Waals equation is a modification of the ideal gas law that accounts for the intermolecular forces and the finite volume of gas molecules. The ‘a’ constant in the Van der Waals equation represents the magnitude of the attractive forces between gas molecules. When the attractive forces are significant, the pressure of the gas will be lower than predicted by the ideal gas law.
Furthermore, the non-zero value of the Van der Waals constant ‘b’ also contributes to the deviation. The ‘b’ constant represents the volume excluded by the gas molecules themselves. When the volume of the gas molecules is significant, the overall volume of the gas will be lower than predicted by the ideal gas law.
In conclusion, real gases deviate from ideal behavior due to various factors, including intermolecular forces, finite volume of gas molecules, complex shapes of molecules, and the non-zero values of the Van der Waals constants ‘a’ and ‘b’. Understanding these deviations is essential for accurate prediction and analysis of real gas behavior in various applications. By accounting for these factors, engineers and scientists can design more efficient and effective systems that operate under real gas conditions.
