Photopilot From Canada, joined Jul 2002, 3142 posts, RR: 15
Reply 3, posted (12 years 2 months 4 weeks 1 day 14 hours ago) and read 3310 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.
N6376m From , joined Dec 1969, posts, RR:
Reply 5, posted (12 years 2 months 4 weeks 18 hours ago) and read 3285 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.
MD11Engineer From Germany, joined Oct 2003, 14968 posts, RR: 61
Reply 6, posted (12 years 2 months 4 weeks 17 hours ago) and read 3273 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.