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Physic Question-can Of Compressed Air  
User currently offlineDavid b. From United States of America, joined Jun 2001, 3148 posts, RR: 5
Posted (12 years 3 months 3 weeks 5 days 9 hours ago) and read 3359 times:

How does a metal can containing compressed high pressure gas get so cold while it is being released? Can someone explain how that happens?

Teenage-know-it-alls should be shot on sight
7 replies: All unread, jump to last
User currently offlineN6376m From , joined Dec 1969, posts, RR:
Reply 1, posted (12 years 3 months 3 weeks 5 days 9 hours ago) and read 3356 times:

Dusting off the old memory banks, I think the answer is found in Boyle's Law Pv=nRT

There is direct relationship between pressure (p) and temperature (t). As the pressure decreases so do the temperature, all other factors remaining constant.

See -



User currently offlineSlamClick From United States of America, joined Nov 2003, 10062 posts, RR: 66
Reply 2, posted (12 years 3 months 3 weeks 5 days 9 hours ago) and read 3351 times:

Correct on Boyle's law and the reciprocal is why compressed gases get hotter.

Aviation-related examples: The adiabatic temperature lapse rate. How an air cycle machine turns hot (from compression) bleed air into cold conditioned air (through expansion) for the cabin.

Happiness is not seeing another trite Ste. Maarten photo all week long.
User currently offlinePhotopilot From Canada, joined Jul 2002, 3163 posts, RR: 15
Reply 3, posted (12 years 3 months 3 weeks 5 days 3 hours ago) and read 3322 times:

Well your partly right and partly wrong with Boyle's Law.

Remember, the can of compressed "gas" that you have and use to blow dust off your equipment is actually a chemical gas. It used to be dichlorodifloromethane which was a florocarbon, however I am not sure what the newer non FL gases are.

This gas is held in a liquid state (shake the can, the liquid sloshes) because the vapor pressure of the gas above the liquid in the can is sufficient to stop the liquid from boiling. Note that this gas boils at a very low temperature, unlike water which boils at 100C.

When you spray the can, you release the pressure above the liquified gas and it starts to boil. If you remember your physics, it takes a tremendous amount of heat to change the liquid to a gaseous state.

This heat comes from the walls of the can which is why they cool down very quickly. If you are holding the can, this heat (energy required to change the state of the liquid to a gas) also is taken from your hand.

Therefore it is not only the loss of pressure as per Boyle that reduces the temperature, but it is also the change of state that requires heat energy.


User currently offlineAirways1 From United Kingdom, joined Jul 1999, 563 posts, RR: 0
Reply 4, posted (12 years 3 months 3 weeks 5 days 3 hours ago) and read 3319 times:

Actually, Boyle's law has nothing to do with it at all.

User currently offlineN6376m From , joined Dec 1969, posts, RR:
Reply 5, posted (12 years 3 months 3 weeks 4 days 8 hours ago) and read 3297 times:

Um not quite right photopilot. The heat that caused the change in state from liquid to gas occurs outside the can. Here's a simple way to prove it.

Invert the can and spray. What will come out is liquid and remains so for less than a second until the spray absorbs enough heat from the evironment to go into a gaseous state. The can will still cool as the pressure in it drops because of the release of the liquid. The pressure inside the can falls why it's volume stays the same PV = nRT.

Furthermore, water boils at 100C only at sea level (or more correctly at the pressure of sea level). If you go to Denver, water boils at a much lower temperature, and therefore that is why you often see "high altitude" cooking instructions on foods such as cakes.

And upon further review, Airways1 does correctly point out the PV=nRT formula is not Boyle's Law but the Ideal Gas Law- though he does do it in sort of a smug manner that reminds me of the dork in high school who always got his ass kicked.

User currently offlineMD11Engineer From Germany, joined Oct 2003, 14968 posts, RR: 61
Reply 6, posted (12 years 3 months 3 weeks 4 days 7 hours ago) and read 3285 times:

What you´ve got to deal with if you expand a gas rapidly is the Joule-Thompson effect. Due to the expansion the gas looses energy. If you have a slow expansion, heat from outside will come in and equalize the amount of temperature lost. In case of a rapid expansion the heat flow from the ambient enviroment isn´t fast enough, this is called an adiabatic expansion. The energy can only come out of the gat itself, causing it to cool down. This effect is widely used in cryogenic machines, like a refrigerator, or in liquifying gases.


Je Suis Charlie et je suis Ahmet aussi
User currently offline777236ER From , joined Dec 1969, posts, RR:
Reply 7, posted (12 years 3 months 3 weeks 3 days 17 hours ago) and read 3249 times:

The answers above just go to show why everyone hates thermodynamics.

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