Assumptions of Ideal Gases: The molecules of an ideal gas are very small, so their volume is assumed zero. There is so much space between the molecules of a gas (as gases have little density) that the only collisions are with the walls of the container and there are no intermolecular forces between the molecules.

When would these assumptions fail? If the volume of the gases was noticeable. If the molecules have intermolecular attractions between them. When would these failures occur? Very low temperature Very high pressure WHY?

Even at standard conditions, some gases deviate noticeably if carefully measured. Why would H 2 O deviate more than H 2 S? Why would CH 3 Cl deviate more than CH 4 ? Why would I 2 deviate more than F 2 ?

For a nonideal gas, PV=nRT does not work, so van der Waals created: (P + n2an2a )(V-nb) = nRT V2V2

“a” is a constant unique for every different gas, but in general it increases as the intermolecular forces increase “n” is involved because the more molecules there are the more intermolecular forces there are P real < P ideal, because the added term is subtracted when solving for P (P + n2an2a )(V-nb) = nRT V2V2

“b” is a constant unique for every different gas, but in general it increases as the volume (size) of the molecules increase “n” is involved because the more molecules there are the more volume they take up V real > V ideal because the subtracted term is added when solving for V (P + n2an2a )(V-nb) = nRT V2V2

Typically nonideal behavior will only change the behavior of a gas about 5% from its ideal behavior, so “a” and “b” are really small values (P + n2an2a )(V-nb) = nRT V2V2

Review for test: Chapter 6 problems: 31, 35, 50, 51, 52, 54, 56, 62, 77, 84, 88, 111 Know nonideal gas issues! Know how to find mole fraction Know reaction rules What is the shape of the line on a volume vs time graph? Graham’s law is not on the blue sheets, so understand how it works and which part is the “root-mean-square speed”