The compressibility factor (Z = PV/nRT) for hydrogen gas (H2) is generally greater than 1 (Z > 1) at room temperature and moderate to high pressures. This behavior is primarily due to:
Correct: B
The compressibility factor Z indicates deviation from ideal behavior. For most gases, at moderate pressures, attractive forces dominate, leading to Z < 1. However, for very light gases like H2 and He, their intermolecular attractive forces (Van der Waals 'a' parameter) are extremely weak. At room temperature, the kinetic energy of these molecules is high enough to largely overcome these weak attractive forces. In such cases, the finite volume occupied by the gas molecules themselves (Van der Waals 'b' parameter) becomes the dominant factor causing deviation. The excluded volume effect leads to a higher effective pressure or volume, making the actual volume (V) larger than the ideal volume, thus Z = PV/nRT > 1. This means significant repulsive forces arise when molecules are pushed close together due to their finite size. Therefore, option B is the primary reason for Z > 1 for H2 at these conditions.