The perceived forces between chlorine gas molecules that contribute to non ideal behavior are most strongly influenced by

Non-ideality in gases Van der Waals equation: P = nRT / (V-nb) - (an²/V²) Two factors cause gases to not behave as ideal gases. 1. gases have finite volumes and take up some space in the container, thus reducing the volume in which they are free to move. This is corrected for by the "b" constant in the van der Waals equation. 2. gas molecules are attracted to each other by intermolecular forces (hydrogen bonding and the three van der Waals forces*). This is corrected for by the "a" constant. Chlorine molecules not polar and are attracted to each other by London dispersion forces. Cl2 is relatively large, with a total of 68 electrons. The strength of London dispersion forces is proportional to the polarizability of the molecule, which in turn is influenced by the total number of electrons and the area over which they are spread. (Any connection to the molar mass is strictly coincidental.) Pressure is always a dependent variable. P depends on V, T and n, plus the corrections in the van der Waals equation. Deviations from ideality (deviations from ideal P) occur more when the volume is small (need "b" correction for the finite molecular volume) or the temperature is low (need "a" correction for the attraction between molecules). * There are three van der Waals forces: Keesom forces (dipole-dipole attraction), Debye forces (induced attraction) and London dispersion forces which all molecules exhibit because of quantum effects.

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