How does the ideal gas differ from real gases ? Give detail explanation with formula

Answer 1

The ideal gas law is a theoretical equation that describes the behavior of ideal gases under certain conditions. An ideal gas is a theoretical concept that assumes the gas molecules have no volume and do not interact with each other. In reality, most gases deviate from ideal behavior due to factors such as the volume of gas molecules and intermolecular forces.

The ideal gas law is represented by the equation:

PV = nRT

where:
P = pressure
V = volume
n = number of moles of gas
R = ideal gas constant
T = temperature

Real gases deviate from ideal behavior due to the following reasons:

1. Volume of gas molecules: In reality, gas molecules have a finite volume which takes up space. This results in the gas molecules being closer together than the ideal gas law predicts, leading to deviations from ideal behavior.

2. Intermolecular forces: Real gases have intermolecular forces that cause them to deviate from ideal behavior. These forces can be attractive or repulsive and affect the pressure, volume, and temperature of the gas.

3. High pressure and low temperature: At high pressures and low temperatures, gases are more likely to deviate from ideal behavior. This is because the gas molecules are more likely to interact with each other and occupy a larger volume than predicted by the ideal gas law.

To account for these deviations, real gases can be described using the Van der Waals equation, which includes correction factors for the factors mentioned above:

(P + a(n/V)^2)(V - nb) = nRT

where:
a = correction factor for intermolecular forces
b = correction factor for volume of gas molecules

In conclusion, while ideal gases follow the ideal gas law under specific conditions, real gases deviate from ideal behavior due to factors such as the volume of gas molecules and intermolecular forces. The Van der Waals equation provides a more accurate description of the behavior of real gases.