Ideal gas Pressure

[premedicine555]

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Hey guys,

1)So I've been at this concept for more than an hour, and I did search the threads but I'm really confused still 🙁. I know that gases that have low pressure, high T are more ideal per se, BUT on TBR page 10 (gas law section) it says that "P(ideal)=Pobserved+[equation]). Why is Pideal greater than the observed pressure??? I understand why the observed pressure is low (because IMF clump it together, less force on wall).... Doesn't that mean that this observed gas is ideal, though, because it has a low pressure? On page 8 it says "low pressure - interacting minimally with one another" for ideal gas criteria... whaaa?

2) Also, is Pideal>Preal; and Vreal>Videal?? I've looked at so many threads, each saying DIFFERENT things cuz apparently EK and TBR/TPR say opposing things.

 

[MedPR]

Vreal>Videal. This one is simple. Real particles have volume, ideals have 0 volume.
Preal<Pideal. This one is a little more complicated. Preal is less because, as you say, there are intermolecular forces and the particles aren't pushing on wall of the container (flask, balloon, whatever). Since P=F/A, less force over the same area = less pressure.

P(ideal)=Pobserved+[equation] because Pideal is greater than P real. Say you calculate P to be 5atm (using P=nRT/V). Since one of the assumptions you made in that calculation (by simply using P=nRT/V) is that the gas particles have no intermolecular forces. So, your calculated P (Pideal) is greater than it really is (Pobserved). To balance that difference out, you have to add a term to the Pobserved side since it is smaller.

The other part you are asking about is the conditions under which a real gas most closely resembles an ideal gas, since ideal gases do not exist. The ideal gas law ignores the volume of gas particles and the intermolecular forces between them when they collide. So, why do low pressure and high temperature make gases behave more ideally than any other condition?

The actual volume occupied by a gas particle becomes more and more insignificant as the volume of your container increases. This is because the gas particles have so much space to spread out and the distances between individual particles is larger than their individual size. From the ideal gas law (and probably intuition), you know that decreasing pressure results in increasing volume. Thus, low pressure for close-to-ideal conditions.

High temperature is for intermolecular attraction/repulsion in collisions. When the gases collide at high temperature (with high kinetic energy), they are less likely to stick to each other and aggregate.

You might also think about it this way:

What happens when you compress (increase pressure) and cool (decrease temperature) a gas? You probably are converting it to a liquid. What happens when you further cool and compress a liquid? Probably going to convert it to a solid. Which phase (solid, liquid, or gas) is most like an ideal gas? Obviously a gas. Think about the phase diagram (the one with a critical point, triple point, etc). Under what conditions is any given point on that diagram the furthest from the liquid and solid phase? Low pressure and high temperature.

 

[premedicine555]

Thank you, that makes sense. So you can have a real gas that has LOWER pressure than an ideal gas, although ideal gases favor lower pressure? I guess i'm getting too caught up in the definitions. I understand your examples, but the whole "gases are closest to ideal at .. low pressure" spiel on page 8 is messing me up. I've also read that increasing pressure by A LOT actually makes volume significant than predicted (ideal).

 

[TXKnight]

Thank you, that makes sense. So you can have a real gas that has LOWER pressure than an ideal gas, although ideal gases favor lower pressure? I guess i'm getting too caught up in the definitions. I understand your examples, but the whole "gases are closest to ideal at .. low pressure" spiel on page 8 is messing me up. I've also read that increasing pressure by A LOT actually makes volume significant than predicted (ideal).

Maybe you are getting caught up in the semantics? ideal gases or real gases do not favor one thing or the other. They are what they are at a given PVT. For the same conditions, then you can compare them. Real gas behavior is closer to ideal (more or less depending on the gas) when at low T and P because of the lower KE, interactions b/w molecules and mean free path. These graphs are useful in understanding that:
http://www.chem.ufl.edu/~itl/4411/lectures/lec_e.html

 

[MedPR]

Thank you, that makes sense. So you can have a real gas that has LOWER pressure than an ideal gas, although ideal gases favor lower pressure? I guess i'm getting too caught up in the definitions. I understand your examples, but the whole "gases are closest to ideal at .. low pressure" spiel on page 8 is messing me up. I've also read that increasing pressure by A LOT actually makes volume significant than predicted (ideal).

Yea, you're just getting confused about the wording. Try this for your thought process.

Which phase is most similar to an ideal gas? A solid? A liquid? A gas?

Obviously any sort of gas (real, ideal, crazy made up crap on an mcat passage) is more closely related to an ideal gas than any sort of solid or any sort of liquid. Now, you should be comfortable with the simple fact that a change in pressure, volume, or temperature will affect the properties of any phase, gases included. Since we want ideal gas, we need to do everything we can to keep our substance in the gas phase. In other words, we need to prevent it from condensing or subliming. How do we accomplish both of those? We put the substance in pressure and temperature conditions that make it impossible for a phase change to occur. In other words, low pressure and high temperature.

Edit: Sorry, not subliming, but depositing (undergoing deposition to a solid). Not sure if depositing is the right term, but deposition is 🙂