Final answer:
The actual gas volume is greater than the predicted ideal gas law volume due to the ideal gas law not accounting for attractions between molecules, which becomes significant at higher pressures.
Step-by-step explanation:
When the actual gas volume is greater than the volume predicted by the ideal gas law, this discrepancy is often due to the ideal gas law not taking into account factors like particle interactions. Specifically, in this case, it is because the ideal gas law does not include a factor for attractions between molecules (c). No significant volume is attributed to the particles in an ideal gas, and no intermolecular forces are considered.
However, in reality, gas molecules do have a finite volume, and they experience forces of attraction with one another. These become significant at higher pressures, meaning the ideal gas law's prediction will differ from the actual volume, as the molecules are physically taking up more space and attracting each other, resulting in less available volume.
At high pressures and lower temperatures, corrections for molecular volume and attractions are required. The van der Waals equation is a modified version of the ideal gas law that can be used to account for these non-ideal behaviors of gases under such conditions.