In which scenario would gas behavior closely approximate ideal conditions?

Gas behavior closely approximates ideal conditions at high temperatures and low pressures. This is due to the fact that, under these conditions, gas particles have greater kinetic energy, which causes them to move more rapidly and occupy a larger volume compared to the space between them, thus reducing the interactions between particles. Ideal gas behavior assumes that the volume of gas particles themselves is negligible, and that there are no intermolecular forces acting on them; these assumptions hold true more effectively at high temperatures where kinetic energy overcomes any attractions or repulsions. Furthermore, at low pressures, gas molecules are more spread out, further decreasing the likelihood of interactions.

Other scenarios, such as low temperatures and high pressures, tend to deviate from ideal gas behavior because particles are closer together, which increases the effects of intermolecular forces and the volume of the particles themselves becomes significant. Moderate pressures and temperatures may not consistently reflect ideal behavior as well, leading to deviations based on specific gas interactions. Standard conditions might suggest ideal behavior but do not guarantee it in all cases, especially when considering different types of gases with varying molecular interactions.