Ideal gases and real gases differ in their behavior under certain conditions. Ideal gases are theoretical gases that obey the ideal gas law, while real gases do not always follow this law due to factors such as molecular volume and intermolecular forces.
The ideal gas law is represented by the equation:
PV = nRT
Where:
P = pressure
V = volume
n = number of moles
R = gas constant
T = temperature
Ideal gases assume that the gas particles have zero volume and do not interact with each other, which leads to the simplified ideal gas law. However, in reality, gas particles do have volume and do interact with each other, leading to deviations from the ideal gas law.
Real gases can be modeled using the van der Waals equation, which accounts for the volume of the gas particles and the attractive forces between them:
(P + a(n/V)^2)(V - nb) = nRT
Where:
a = constant related to molecular attraction
b = constant related to molecular volume
The van der Waals equation helps correct for the deviations in behavior of real gases from ideal gas behavior. For example, at high pressures or low temperatures, real gases deviate more from ideal gas behavior due to the closer proximity of particles and the increased importance of intermolecular forces.
In summary, ideal gases and real gases differ in their behavior due to the assumptions made in the ideal gas law and the corrections needed in the van der Waals equation to account for real gas behavior. Ideal gases are theoretical and follow the ideal gas law, while real gases have volume and interact with each other, leading to deviations from this law.
How does ideal gas differ from real gas ? Give detail explain with formula and inicials meaning with examples
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