(A) : At high pressure, real volume is less than ideal volume
(R) : At high pressure gas molecules cannot strike the walls of the container with full impact and experience a net backward pull

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(A) is true but (R) is false

Both (A) and (R) are correct and (R) is the correct explanation of (A)

Both (A) and (R) are false

Both (A) and (R) are correct but (R) is not the correct explanation of (A)

answer is B.

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Detailed Solution

At high pressures, the volume occupied by the gas molecules becomes significant compared to the total volume of the gas, which leads to deviations from ideal behavior.
In the ideal gas law, the volume of the gas molecules is assumed to be negligible compared to the total volume of the gas.
However, in real gases, the volume of the gas molecules is not negligible and cannot be ignored, especially at high pressures.
As the pressure increases, the volume occupied by the gas molecules becomes more significant, causing the real volume of the gas to be less than the ideal volume predicted by the ideal gas law.
This effect is more pronounced for gases with larger molecular sizes or stronger intermolecular forces, as these factors increase the volume of the gas molecules and make the deviations from ideal behavior more significant.
At high pressures, gas molecules are more closely packed together and have less space to move around.
As a result, the molecules cannot strike the walls of the container with full impact, and they experience a net backward pull due to intermolecular forces.
This backward pull reduces the pressure exerted by the gas, which also leads to a deviation from ideal behavior.
In addition, at high pressures, the average kinetic energy of gas molecules increases, which can lead to more frequent and more energetic collisions between the molecules.
This can also lead to an increase in the net backward pull experienced by the molecules, further reducing the pressure exerted by the gas and increasing the deviation from ideal behavior.

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