Question:
Ideal gas and real gas have major difference of
Ideal gas and real gas have major difference of
Updated On: Jun 20, 2022
- phase transition
- temperature
- pressure
- None of the above
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The Correct Option is C
Solution and Explanation
No attractive or repulsive force acts between gas molecules in an ideal gas whereas for a real gas due to intermolecular force in real gases. Molecules do not exert that force on the wall which they would have exerted in the absence of intermolecular force. Therefore, the observed the observed pressure p of the gas will be less than that present in the absence of intermolecular force. Hence, the attractive becomes $ \bigg( p + \frac{a}{V_2}\bigg)$.
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Concepts Used:
Kinetic Molecular Theory of Gases
Postulates of Kinetic Theory of Gases:
- Gases consist of particles in constant, random motion. They continue in a straight line until they collide with each other or the walls of their container.
- Particles are point masses with no volume. The particles are so small compared to the space between them, that we do not consider their size in ideal gases.
- Gas pressure is due to the molecules colliding with the walls of the container. All of these collisions are perfectly elastic, meaning that there is no change in energy of either the particles or the wall upon collision. No energy is lost or gained from collisions. The time it takes to collide is negligible compared with the time between collisions.
- The kinetic energy of a gas is a measure of its Kelvin temperature. Individual gas molecules have different speeds, but the temperature and
kinetic energy of the gas refer to the average of these speeds. - The average kinetic energy of a gas particle is directly proportional to the temperature. An increase in temperature increases the speed in which the gas molecules move.
- All gases at a given temperature have the same average kinetic energy.
- Lighter gas molecules move faster than heavier molecules.
