Why the measured pressure of a gas under conditions that are very close to those that would result in condensation will be lower than what the ideal gas law would predict?

Inter-molecular forces and molecular volumes are the chief reasons for lower measured pressure

Explanation:

The kinetic theory assumes that gas particles occupy a negligible fraction of the total volume of the gas. It also assumes that the force of attraction between gas molecules is zero.

However, during high pressure, the volume of the gas particles are not negligible compare to the total gas volume and as such the volume of a real gas under such condition is higher than the Ideal gas. Vander-waal attempted to modify the ideal gas equation by subtracting the excess volume from the ideal equation. The increased volume is the reason the measured pressure of a real gas is less than an ideal gas

On the other hand, close to condensation, the other assumption of negligible forces of attraction becomes invalid. As inter-molecular distances decrease, inter-molecular forces increase reducing the bombardment of the wall of the container due to restricted particle movement and lower measured gas pressure.