Final answer:
A gas behaves more like an ideal gas at low pressure and high temperature, where molecular volume and intermolecular forces are negligible. At high pressure and low temperature, gases deviate from ideal behavior due to significant molecular interactions and reduced space between molecules.
Step-by-step explanation:
Gases behave more like ideal gases at certain temperatures and pressures. Specifically, low pressure and high temperature are the conditions that align more closely with the assumptions made by the ideal gas law. This is because at low pressure, gas molecules are well-separated, minimizing the intermolecular forces that can cause deviations from ideal behavior. At high temperature, kinetic energy is greater, which allows gas molecules to overcome any intermolecular attractive forces that might exist, behaving as if they are non-interacting particles.
On the other hand, a gas deviates from ideal behavior at high pressure and low temperature. High pressure reduces the volume between particles, thus the volume of the gas particles cannot be considered negligible. At low temperatures, the overall kinetic energy of the gas decreases, magnifying the effect of intermolecular attractions, which again leads to non-ideal behavior. Indeed, at sufficiently low temperatures and high pressures, gases can liquefy, which is certainly not ideal gas behavior.
The ideal gas law is most accurate in describing a real gas when intermolecular attractions and the volume occupied by the gas molecules are insignificant. This typically occurs at low pressures where molecules are far apart and at higher temperatures where their kinetic energy dominates their behavior, resulting in gas particles acting independently of each other.