I have two questions, so if anyone could answer them it would be very helpful: 1) How can you cool a gas by increasing the volume, even though PV=nRT indicates otherwise. I thought if you increase volume, you also increase temperature? 2) High pressure and low temperature cause deviations from ideal gas behavior, so gases deviate from ideal behavior when pressure is >10atm and temperatures near the boiling points. Shouldn't we decrease temperature? From PV=nRT if we increase temperature we can increase pressure (and decrease volume), creating intermolecular forces. But if this is true, gases deviate at extremely high temperature and extremely low temperatures, since gases come together when temperature decreases. please help!

Volume increases, pressure decreases, temperature decreases. If we compress a gas, we do work on the system, increase it's energy and raise it's temperature. I don't know about the second question...I haven't gotten to ideal gases yet in my content review.

I think you are confusing a few concepts here. First, ask yourself the question: Why do gases deviate from ideal gas behavior at high pressure and low temperature? The reason is that since the molecules are closer together, the intermolecular forces become more significant. At high temperature, gases have lots of kinetic energy, so the intermolecular forces are less significant. The same happens at low pressure when the molecules are spaced farther apart.

One thing you might want to do is understand the concepts around this equation. One cannot just look at the equation and derive concepts from it because of the confusion it stems. Changing the volume would cause the PRESSURE to decrease which would cause the temperature to decrease. If you add heat to the system, you can cause the pressure to increase and the volume to increase. (the pressure increase is only to equilibreate to the original volume) There is the pressure variable in there too. And to your second question, the gases have such high kinetic energy that they overcome the intermolecular forces that would normally occur at low temperatures. Hope that helps.

perhaps you should re-read your q, bc it doesn't seem to entirely make sense. based on what i think you're asking... consider the assumptions behind the ideal gas equation - that molecules occupy no volume and that intermolecular interactions do not occur. the ideal gas equation is always wrong because these assumptions are never entirely true. they become more significant, however, at high pressure (first assumption becomes worse) and lower temperature (second assumption becomes worse). i don't understand "shouldn't we decrease the temperature?"...

Both questions can be answered by telling you that the ideal gas equation can be used the way you're using it if, and only if, two parameters are variable. Ok, so PV=nRT, so you usually think, oh, ok, i'll increase the pressure so I'll increase the temperature. the problem is, that's only true if volume doesn't change as a result of that pressure increase. if it does, then the temperature change will be a function of both pressure change and volume change. so, for your first question, which is describing adiabatic expansion of a gas, the reason you get cooling of a gas (and trust me, I used a CO2 cylinder without a regulator, ice crystals started forming around the valve because of the cooling) must be derived: dE=dq+dw (or you may have learned dq-dw, the sign of dw is arbitrary, just be consistent). So, dE is defined as CvdT. For an adiabatic process, dq is always 0. dw=-PdV (the sign is negative because work is done by the system as volume/dV increases so the system is losing energy) NOW you can apply PV=nRT: Rearrange so you get P=nRT/V Plug into dw=-PdV to get dw=-nRT/VdV so now you just have: CvdT=-nRT/VdV rearrange to get T on the other side: (Cv/T)dT=-(nR/V)dV now integrate both sides: (n and R are constants, and Cv can be taken to be a constant in many cases) sooo Int[(Cv/T)dT]=Int[(-nR/V)dV] take constants out: CvInt[dT/T]=-nRInt[dV/V] we're integrating over T1 to T2 and V1 to V2 so we get: Cv*Ln[T2/T1]=-nR*Ln[V2/V1] because the sign is negative, you flip the Ln Cv*Ln[T2/T1]=nR*Ln[V1/V2] sooooo as you can see as you increase temperature, (so T2/T1 is greater than 1), V1/V2 is also greater than 1 which means that you're decreasing volume. the converse is also true. For your second question, think of it like this: high temperature, if all else is constant, will tend to want to get molecules as far apart as possible. low pressure, if all else is constant, will want to get molecules as far apart as possible. if you have both high temperatures and low pressure, you'll be getting the molecules as far apart as possible. if you increase pressure, you're squeezing the molecules together. if you decrease temperature, molecules have lower avg kinetic energy so they tend to get closer together, again, if all else is constant. that's why high temp, low pressure makes it most ideal: it gets everything far apart to minimize molecular interactions. hope this helps

There are plenty of threads about this. I know I've replied to a handful. Search and you'll find some more answers.

lol. I was actually going to say the exact same thing. I know this like the back of my hand thanks to the past couple months here.