Real gases volume changes

[MrNeuro]

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Can someone explain to me why the volume is a little more when pressure is increased? and why the pressure is a little smaller when the pressure is halved???

i get the whole empty space thing but when it comes to total volume thats when i get lost....

so assume you have 1 mol of gast at STP

if you double the pressure the volumes going to be a little more than 11.2 L and
if you half the pressure the volumes going to be a little less than 22.4 L

how in the world does that work out?

the gas is taking up some volume so how is it giving a little less volume than 22.4 shouldn't it be greater

Videal = V container - nb

Vcontainer = Videal + nb

the part thats even more mind bending is the fact that Van der Waals agrees....

where am i going wrong?

 

[Dasypus]

You have to remember what it is we're idealizing away when we talk about 'ideal gasses'. In an ideal gas, the individual particles of gas take up no space (we model them as points) and have no intermolecular interactions.

Now, no real gas is like this, but within a broad range it's a fair approximation. The approximation breaks down when we put a lot of real gaseous particles in a (relatively) small space. Try to squish them down really small, say by lowering a tight-fitting piston on a cylinder, and they won't squish as small as an ideal gas would because the particles themselves take up room in that space, the way ideal point particles don't.

On the other hand, if you give them more space, say by raising the tight-fitting piston, you won't have to raise it as far as you thought, because now the interesting deviation from ideal isn't that the particles take up room, but rather that they interact with each other. Everything interacts to some extent; that's the van der Waals (aka London dispersion) forces at work at the very least --there may be stronger ones too, though usually if you've got stronger attractive forces, you've got a liquid. So the particles are holding together a wee bit more than 'ideal', and these attractions can work over a fair range of pressures because everying's tumbling around in a gas. Thus at lower pressures this effect is stronger than the fact that they take up room, because that makes the big difference when they're squished down really tight.

Hope that helps!

 

[chiddler]

halving pressure can be thought of as halving the distance between gas particles.

try to look at it again with this knowledge.

 

[Dasypus]

halving pressure can be thought of as halving the distance between gas particles.

try to look at it again with this knowledge.

I'm pretty sure that's backwards. Up the pressure, force the distances to be smaller. Lower the pressure, particles have more space and spread out.

 

[chiddler]

oops sorry.

 

[Dasypus]

oops sorry.

Edit and change it? 😉

 

[Dasypus]

Here, I don't know if this will help or hinder, but here's a graph of the real behavior of N2 under pressure: http://www.chemguide.co.uk/physical/kt/realgraphs1.jpg

If it were an ideal gas, all of those curves would be straight lines at y=1. The bit where it drops below 1 is where intermolecular forces are pulling together, so the actual volume is less than expected. It crosses 1 as that effect is cancelled out by the volume of the gas particles. That ratio continues going up as the pressure does because at higher pressures more of the available space is filled with particles, so the real volume keeps getting bigger than you'd expect it to be.