04-12-2018, 03:32 PM

Hello all,

I have to teach the following problem this weekend and I am having trouble with one part of it. Any help would be greatly appreciated!!!!

2. In a reaction at equilibrium, it is concluded that helium gas behaves more ideally than carbon dioxide. Which of the following accurately explains why this is so?

I. CO2 exerts a greater pressure because its molecules have lesser volume.

II. The intermolecular forces between CO2 molecules are stronger than those in He.

III. Helium molecules have greater kinetic energy, and therefore behave more ideally.

A. I only

B. II only

C. II and III only

D. I, II, and III

First, I am not sure what "reaction" would allow one to discern this, but that is besides the point.

My thoughts:

I. Wrong because a molecule of CO2(g) would occupy more volume than a molecule of helium(g).

II. True. He is a noble gas and therefore has minimal intermolecular forces. CO2 on the other is more polarizable and will have more temporary dipoles/van der waals forces.

III. I believe this is False, but I cannot come up with a completely exhaustive reason why. For gases I think of kinetic energy as average kinetic energy, and therefore equal to 3/2RT. No temperature is implied in this problem. I assume if the temperature of the 2 different gases was the same then they would have the same Kinetic energy (if we can assume they are both ideal, which seems unlikely in this problem). I have always thought of temperature as relying on the conditions that the gas is subject to rather than an intrinsic property of the gas itself. Also Van Der Waals equation does not directly account for variations in Temperature: (P+a(n/V)^2) X (V - nb) = nRT.

So I thought about what would happen if you took two 1L containers, filled one with one mol CO2 and the other with one mol of He. Hold all conditions constant (except T), and then measured the temperature of each gas? Thinking about deviations from ideality, I noted that since CO2 has more attractive forces and more volume (on a molecular basis), these factors would lead to an decreased pressure and slightly increased volume. If you plug those into PV/nR = T then I suppose there would be a net effect of increasing the temperature and thus the KE relative to the container of He.

I am pretty sure that my thinking is flawed-- could someone please help me find where I have gone wrong?

Thank you so much!!!!

I have to teach the following problem this weekend and I am having trouble with one part of it. Any help would be greatly appreciated!!!!

2. In a reaction at equilibrium, it is concluded that helium gas behaves more ideally than carbon dioxide. Which of the following accurately explains why this is so?

I. CO2 exerts a greater pressure because its molecules have lesser volume.

II. The intermolecular forces between CO2 molecules are stronger than those in He.

III. Helium molecules have greater kinetic energy, and therefore behave more ideally.

A. I only

B. II only

C. II and III only

D. I, II, and III

First, I am not sure what "reaction" would allow one to discern this, but that is besides the point.

My thoughts:

I. Wrong because a molecule of CO2(g) would occupy more volume than a molecule of helium(g).

II. True. He is a noble gas and therefore has minimal intermolecular forces. CO2 on the other is more polarizable and will have more temporary dipoles/van der waals forces.

III. I believe this is False, but I cannot come up with a completely exhaustive reason why. For gases I think of kinetic energy as average kinetic energy, and therefore equal to 3/2RT. No temperature is implied in this problem. I assume if the temperature of the 2 different gases was the same then they would have the same Kinetic energy (if we can assume they are both ideal, which seems unlikely in this problem). I have always thought of temperature as relying on the conditions that the gas is subject to rather than an intrinsic property of the gas itself. Also Van Der Waals equation does not directly account for variations in Temperature: (P+a(n/V)^2) X (V - nb) = nRT.

So I thought about what would happen if you took two 1L containers, filled one with one mol CO2 and the other with one mol of He. Hold all conditions constant (except T), and then measured the temperature of each gas? Thinking about deviations from ideality, I noted that since CO2 has more attractive forces and more volume (on a molecular basis), these factors would lead to an decreased pressure and slightly increased volume. If you plug those into PV/nR = T then I suppose there would be a net effect of increasing the temperature and thus the KE relative to the container of He.

I am pretty sure that my thinking is flawed-- could someone please help me find where I have gone wrong?

Thank you so much!!!!