Final answer:
Real gases behave most like an ideal gas under high temperatures and low pressures. At high pressures and low temperatures, real gases deviate significantly from ideal behavior due to significant intermolecular forces and the non-negligible volume of the gas molecules.
Step-by-step explanation:
To determine which gas is most likely to resemble the ideal gas at higher pressures, we have to look at conditions under which real gases approximate ideal gas behavior. Generally, real gases behave more like ideal gases at high temperatures and low pressures. The factors causing deviation from ideal behavior include intermolecular forces and the finite volume of gas molecules. These effects are less noticeable at high temperatures, as the increased kinetic energy of the molecules tends to overcome the intermolecular forces, and less noticeable at low pressures, where the volume occupied by gas molecules is negligible compared to the volume of the container.
At high pressures, however, these assumptions break down. The volume of the gas molecules becomes significant, and intermolecular forces can no longer be ignored, especially at low temperatures. A real gas is expected to deviate from ideal behavior at high pressures and low temperatures because these conditions don't allow the gas particles to move freely and independently.
The van der Waals equation corrects for these deviations by considering the volume of gas molecules and intermolecular forces. As pressure increases, however, the likelihood of gases behaving ideally decreases. Therefore, no real gas resembles an ideal gas perfectly at high pressures, but they may still come close under appropriate high-temperature conditions.