[mike2niner4 (OP)]

I understand that when a gas expands it cools but why does it cool?

Thanks, Mike  []

[lyner]

Imagine the gas was in a cylinder and expanded by pushing against a piston. The movement would involve work being done. This work would need energy, which would come from the kinetic energy of the gas molecules hitting the piston- reducing their average speed. Taking KE away corresponds to lowering the temperature of the gas.

[Geezer]

As Sophiecentaur says, it's because of the kinetic energy of the gas molecules.

It might also help to turn the question on it's head. As in, why does air heat up when it is compressed? - well known in bicycle pumps.

The air contains a certain number of molecules. The molecules have a certain amount of heat energy (in the form of kinetic energy). When the air is compressed the total heat energy now occupies a smaller volume. Unless heat is rapidly removed while the air is being compressed, the air's temperature has to increase because the heat energy per unit volume is increasing.

The expansion process is the exact opposite, so unless heat is added while the air is expanding, the temperature has to decrease.

The compression heating effect can be quite dramatic. Before the development of matches, air compression devices were used for lighting fires.

[John Chapman]

When the air is compressed the total heat energy now occupies a smaller volume. Unless heat is rapidly removed while the air is being compressed, the air's temperature has to increase because the heat energy per unit volume is increasing.

That's a really good explanation. I've often wondered how that works. What about in diesel engines though? Would the compression ratio be enough to reach the high temperatures needed for combustion using your explanation?

Also, is that effect the same thing that makes an elastic band warm up when you stretch it? It seems sort of the wrong way round. If you cut an elastic band (a chunky one works best) to get a length of elastic then stretch it in front of your face and immediately touch the elastic on your top lip you'll feel that it is noticeably warm. If you let it cool for a few seconds, then relax it and touch your top lip again it is now noticeably cold. Doesn't the elastic band fills the same volume whether it is stretched or relaxed? I'm not sure. But how does that fit into your explanation above?
 

[chris]

Hi John

we did this experiment as a Kitchen Science.

http://www.thenakedscientists.com/HTML/content/kitchenscience/exp/make-a-rubber-fridge/

The bottom line is that when you stretch the rubber band to begin with you are doing work to straighten out the randomly-coiled polymer chains in the rubber. You are therefore adding energy to the rubber so it becomes hotter.

When you let you the process is reversed. Now the rubber does work pulling the straightened chains back  to their initial compact configuration. Since the elastic has done work, using energy, the rubber must lose energy in the process and therefore it becomes colder.

In the link above Dave's put a nice interactive animation that enables you to demonstrate this for yourself.

Chris

[John Chapman]

That's a fun computer model. Clever old Dave.

I'm still not clear about this, though? Are you saying that something will necessarily get warmer if you put energy into it? Will a wrist watch get warmer if you wind it up or a jar of pickles if you lift it on to a higher shelf?

Sorry for highjacking your thread, mike2niner4. But this is a vaguely similar question to the original one, which I think has now been answered by Geezer's very plausible sounding explanation.
 

[lyner]

geezer

the air's temperature has to increase because the heat energy per unit volume is increasing.

But the mass of gas is the same. Unless you say that the specific heat capacity changes, the temperature shouldn't change by that argument.
The argument must hang on work done on or by the gas and the resulting change in internal energy.
This is very standard 'book work' concerning adiabatic changes.

[John Chapman]

Ahhh, I've heard about adiabatic cooling in respect of meteorology. Moving air hits a mountain or hill and is forced upwards. With higher altitude comes lower pressure and the air expands and cools. Cold air holds less water vapour than warm air and at some point up the mountain the air will reach it's dew point (the water saturation point of the air) and the water vapour will condense out as cloud. This is why cartoon mountains are often drawn with a cloud circling it's peak.
 

[Turveyd]

Nice,  when you compress the gas,  the heat energy compresses which creates heat,  which is then lost as this will resume to room temperature,  so when you release the gas your releasing gas with less heat nice.

Never considered that before.

I like science.

[Geezer]

But the mass of gas is the same. Unless you say that the specific heat capacity changes, the temperature shouldn't change by that argument.
The argument must hang on work done on or by the gas and the resulting change in internal energy.
This is very standard 'book work' concerning adiabatic changes.

Sophiecentaur

The question was "why does the gas cool", not "how much does the gas cool". Certainly the work done on or by the gas can lead us to the temperature of the gas at a certain pressure and volume, but it hardly explains why the temperature of the gas changes.

We could also say, well, P times V is proportional to T, because it's a well known law, but that does not provide too much insight either.

The temperature of a gas is proportional to the kinetic energy in the molecules of the gas. If we enclose a certain number of molecules of the gas and reduce the volume of the enclosure (which does require work) there are still the same number of molecules in the enclosure, but they now occupy a smaller space. This increases the kinetic energy of the enclosed gas, and hence the temperature. The energy per unit volume has also increased.

[lyner]

The reason "why" it cools is that the molecules have lost average KE. This is 'because' they have lost it in collisions with the retreating walls of the container* - their recoil speed is less than their approach speed, having transferred some energy to the container by doing work. There may be a deeper 'why' to answer but, in terms of kinetic theory, I can't think of one.

*an actual container is not necessary for this argument to apply - you can replace the walls of the container by molecules of a surrounding gas.

[Turveyd]

It cools when you compress it,  as there is the same heat in a smaller area if you felt a can / oxygen cylinder when it's being compressed it's warm / hot,  so when it leaves it's lost the temperature already.

If you put your can in the oven and heat it to 100c most likely all through then the gas should come out near room temperatue.


Don't try this the can will explode

[Geezer]

It cools when you compress it,  as there is the same heat in a smaller area if you felt a can / oxygen cylinder when it's being compressed it's warm / hot,  so when it leaves it's lost the temperature already.

If you put your can in the oven and heat it to 100c most likely all through then the gas should come out near room temperatue.


Don't try this the can will explode

Turveyd:
Ah - well yes. When you compress gas it heats, and the heat can escape, if you let it. The situation we were debating does not allow the heat to escape. If you look back in the thread you'll notice that Sophiecentaur mentioned adiabatic which is the term used to describe a process where there is no heat transfer.

If there is heat transfer, the point I made about compression and expansion being reversible would not be valid. As you point out, if energy is allowed to escape from the can in the form of heat, you would have to add back the energy that escaped if you wanted the air to return to the original temperature when it expanded again.

Sophiecentaur can provide a more rigorous explanation if required.

[Bored chemist]

Imagine I have some gas in a container. At the moment all the gas is in one side of the container and it's held there by a thin wall. On the other side of the wall (i.e. in the other half of the container) there's a vacuum.
Now imagine I take the thin wall away. The molecules of gas that would have hit it carry on to the other half of the container at exactly the same speed. They hit it and bounce off- just as they would have done from the wall I took away.
After a very short time the molecules are randomly distributed. They occupy twice the volume they did so the pressure is halved.
But no work has been done on or by the gas. It hasn't gained or lost any energy. It is, therefore at exactly the same temperature as it was before.

The temperature of an ideal gas expanding into a vacuum does not change.
You need to be clear about answering questions like "why does the gas cool?"

Sometimes it doesn't.

[lyner]

Correct. But this is a very 'ideal' situation, not often encountered - certainly not naturally - in the weather, for instance.
Most scenarios involve some work being done, though. A piston or a nozzle involve energy loss. A gas which expands by exerting a force times a distance will get  cool.  Then there are Van de Wall's forces which mean that a non ideal gas will still cool down - even when exploding into a vacuum.

[Bored chemist]

If you let high pressure hydrogen or helium out of a cylinder they warm up.

[lyner]

I never knew that. Go on then. Give us an explanation.
Are some special conditions necessary?
The energy must come from somewhere. He is monatomic so how....?

[Nizzle]

I think it warms up because other molecules from the environment bump into the high pressure Hydrogen or Helium (which has very little movement/E_k) and transfer some kinetic energy that way.

Or are we still talking about a vacuum environment here?

[Bored chemist]

There's lots of stuff about it here.
http://en.wikipedia.org/wiki/Joule%E2%80%93Thomson_effect

[lyner]

How dumb of me. I remember being here before.  It's that inversion temperature thing in the Joule Kelvin effect. A very nice argument to explain it.