non-ideal gases

[Deepa100]

compressibility ratio, Z, as a function of pressure, Z = PV/nRT.

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It is expected that as the attractive forces on a molecule increase, the compressibility ratio will:
increase linearly.
decrease linearly.
increase proportional to 1/a.
remain constant.

[Show/hide explanation]

Think about how attractive forces would affect either pressure or volume. If the attractive forces are high, then molecules are "holding on" to each other with more strength. Therefore, the pressure (which is the pressure of the molecules against the walls of the container) would be effectively decreased. A decrease in pressure would lead to a decrease in the compressibility ratio. Choice (B) is correct. Notice that knowing quantitatively how it decreases (i.e., linearly, etc.) is not important to answer this question.

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My Q: If the gas is non-ideal, pressure increases and volume decreases due to intermolecular forces.

So, why does this solution say pressure decreases?

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[Lukkie]

My Q: If the gas is non-ideal, pressure increases and volume decreases due to intermolecular forces.

WRONG. it is the OPPOSITE.

in an ideal gas, the gas molecules are point particles. they are basically a miniscule, almost zero volume. this is not true in real gases, however. gas molecules do have a volume. i think you are getting confused on what tje 'volume' REQUIREMENT means in the ideal gas approximation - its not the overall volume, but rather the volume of a single molecule.

Also, in an ideal gas, the molecules don't exhibit any forces on eachother other than collisions. however, in a real gas, you have attractive forces. now think of a big tin can with millions of gas molecules. they are moving in every direction, and are distributed pretty evenly all over. however, if you had to pick a 'center of mass' for the system, so to speak, it would be right in the middle of the tin can. so you can imagine that all the gas molecules will generally be attracted to the middle. now what is pressure? basically, its the force at which particles hit the side of the tin can. if molecules are being ATTRACTED to the MIDDLE of the tin can, they have less force going outwards. so there is LESS PRESSURE.

 

[andafoo]

compressibility ratio, Z, as a function of pressure, Z = PV/nRT.

==============
It is expected that as the attractive forces on a molecule increase, the compressibility ratio will:
increase linearly.
decrease linearly.
increase proportional to 1/a.
remain constant.

[Show/hide explanation]

Think about how attractive forces would affect either pressure or volume. If the attractive forces are high, then molecules are "holding on" to each other with more strength. Therefore, the pressure (which is the pressure of the molecules against the walls of the container) would be effectively decreased. A decrease in pressure would lead to a decrease in the compressibility ratio. Choice (B) is correct. Notice that knowing quantitatively how it decreases (i.e., linearly, etc.) is not important to answer this question.

========
My Q: If the gas is non-ideal, pressure increases and volume decreases due to intermolecular forces.

So, why does this solution say pressure decreases?

You shouldn't memorize that in non-ideal gases pressure always increases and volume always decreases. Depending on what equation of state you use you can different behavior in real gases.

Using the van der waals equation you can see that Z is directly prop. to P. Seeing as the intermolecular forces in this case are causing the molecules to attract to each other, pressure will be reduced. Make sense?

 

[Deepa100]

You shouldn't memorize that in non-ideal gases pressure always increases and volume always decreases. Depending on what equation of state you use you can different behavior in real gases.

Using the van der waals equation you can see that Z is directly prop. to P. Seeing as the intermolecular forces in this case are causing the molecules to attract to each other, pressure will be reduced. Make sense?

The equation for the non-ideal gases is also given in the passage. Cut and paste did not work... It basically shows (P+X)(V-Y) = nRT
where the term X has a constant a that is dependent on the intermolecular forces. So, I was not memorizing it, I was just using the passage info.

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Here is a cut and paste of the entire passage
he hypothetical compounds known as ideal gases often give chemists great insight into the behavior of real gases. The ideal gas approximation arises from basic trends deduced in the behavior of real gases at low pressure. Boyle's law, for instance, results directly from the assumptions underlying the kinetic theory of gases:

1. A pure gas consists of a large quantity of identical molecules separated by very large distances relative to the size of the molecules.
2. Molecules move in random directions with a distribution of speeds.
3. Molecules exert no forces on one another between collisions.
4. Collisions of molecules are elastic; no energy is lost.

Though the ideal gas approximation is accurate near atmospheric pressure, at high densities the behavior of real gases diverges from the ideal model, since these conditions invalidate one or more of the kinetic theory's assumptions. To account for these deviations, the nineteenth-century Dutch physicist Johannes van der Waals proposed appropriate corrections to the ideal gas law, expressed by the van der Waals equation of state, Equation 1.

Equation 1
XXXXXXXXMissing equation goes hereXXXXXXXXXXXXXXXXX
The parameters a and b are functions of the physical characteristics of specific molecules. The van der Waals constants for N2 and H2 are listed below in Table 1.

Table 1 Van der Waals Constants for N2 and H2

When deciding whether to treat an unknown gas as ideal, chemists must determine how accurate they require their results to be, and include multiple factors in their decision: ambient pressure, intermolecular attractions, and the size of gas molecules.

One such chemist attempts to analyze an unknown compound twice. The chemist assumes the gas is ideal in his first analysis and real in his second. To gauge the accuracy of his first assumption, the chemist calculates the compressibility ratio, Z, as a function of pressure, Z = PV/nRT.