Final answer:
A real gas differs from an ideal gas because its molecules have some volume and experience some attraction for each other, affecting its compressibility and behavior compared to the ideal gas, especially at high pressures. So the correct option is B.
Step-by-step explanation:
A real gas differs from an ideal gas in that the molecules of a real gas have some volume and some attraction for each other. This is in contrast to the ideal gas, where the molecules are assumed to have no significant volume and no attractive or repulsive forces between them. At low pressures and high temperatures, these deviations from the ideal behavior are generally negligible, and real gases tend to behave similarly to ideal gases. However, at high pressures, the molecules of a real gas are compressed closer together, thereby reducing the amount of space between them. This results in the molecules' volume becoming appreciable relative to the total volume of the gas and leading to less compressibility at high pressures. Furthermore, the attractive forces between molecules in a real gas can lead to behavior that deviates from what is predicted by the kinetic molecular theory of gases and ideal gas laws.