For a perfect ideal gas, the compression factor always equals exactly 1. This happens because the ideal gas law assumes particles have absolutely zero volume. It also assumes the gas particles never attract or repel each other.
Real gases deviate from this perfect behavior because their molecules have actual physical size. Real molecules also experience intermolecular forces that pull them toward one another. At low to moderate pressures, these attractive forces pull the molecules closer together.
This makes the real gas take up less volume than an ideal gas. As a result, the compression factor Z drops slightly below the value of 1. At very high pressures, the physical volume of the molecules becomes much more important.
The molecules get squeezed so close together that they start repelling each other. These strong repulsive forces push the molecules apart and increase the overall volume. This causes the compression factor Z to rise above 1 at high pressures.
Plotting this factor helps chemists understand how different gases handle extreme industrial conditions. A common misconception is that all gases act identically under high pressure. In reality, each specific gas deviates from ideal behavior in its own unique way.