Could I get some help here? This is from a Kaplan flash card:
How do actual volume and predicted volume of a gas compare at:
1)moderately high pressure?
2)Extremely high pressure?
3)very low temp?
Answer:
1) less than predicted
2) more than predicted
3) less than predicted
An ideal gas has no intermolecular interactions, and that is the key to this problem...
1) At moderately high pressure, the molecules are forced to be closer together, which will create some intermolecular interactions. These interactions (read: bonds, albeit only for microsecond intervals) pull the molecules together so the the actual volume will be less than predicted by the ideal gas law, which assumes no interactions.
2) At extremely high pressures, the molecules are very close together and, in this situation, the volume of each individual molecule comes into play. The ideal gas law assumes that each molecule actually has no volume, but as you compress the gas, at some point the fact that the molecules actually do have volume comes into play, and the actual volume will be greater than what is predicted by the ideal gas law. The extreme proximity of the molecules in this case also introduces repulsive forces between the molecules, which further serves to increase the volume of the actual gas. To differentiate between case 1 and 2, keep in mind that molecules will attract when they are close (VanDer Waal's interactions) but will repel when they are TOO close (electrostatic interactions).
3) This is similar to #1. The low temperature means the molecules are moving slower and are more readily able to interact. The interactions give the actual gas a smaller volume than what was predicted by the ideal gas law, which assumes no interactions.
