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WAY off topic (was Re: nitrous in your car?)
"God" wrote:
>True - the pressure will increase. This pressure increase is due, however,
>primarily to the increasing equilibrium vapor pressure of the liquid N2O
What if the tank is nearly empty (i.e. little or no liquid remaining)? This
is the scenario that bottle heaters are designed to address.
>What happens when a sample of gas/liquid at equilibrium is
>compressed (i.e, the volume available is reduced) at constant temperature?
>Some more gas liquifies. The pressure stays essentially constant and does
>not increase as the ideal gas law would predict.
But as V decreases, so does n in this situation. I don't think there's
necessarily a contradiction with the ideal gas law here.
OK, OK, maybe I was a bit hasty in citing PV=nRT...but I still think this
cursory explanation is adequate for the layperson who asked the original
question.
Regards,
Bob W.
-----Original Message-----
From: Robert Myers <rmyers@inetone.net>
To: Robert Wunderlich, DPM <Robert@Wunderlich.com>
Cc: Quattro list <quattro@audifans.com>
Date: Thursday, September 30, 1999 6:40 AM
Subject: Re: nitrous in your car?
>Caution y'all - pedant alert! Delete now unless you want to become
>terminally bored. :-)
>
>At 04:44 AM 09/30/1999 -0500, you wrote:
>>>Once all the liquid has been evaporated then the PV=nRT equation applies.
>>
>>
>>I agree with everything you wrote except the above statement. How can the
>>Ideal gas law NOT apply? Regardless of the state or amount of nitrous in
>>the tank, if you increase the temperature of the tank (without releasing
>>gas) the pressure is going to increase.
>
>True - the pressure will increase. This pressure increase is due, however,
>primarily to the increasing equilibrium vapor pressure of the liquid N2O
>rather than some application of the ideal gas law (which is only a first
>approximation at best). Remember, we are dealing with a change of state
>not just a gas in a close system. This makes the system far more complex
>than just a single phase gaseous system and the ideal gas law as typically
>encountered in a freshman chemistry class is not capable of handling the
>What happens when a sample of gas/liquid at equilibrium is
>compressed (i.e, the volume available is reduced) at constant temperature?
>Some more gas liquifies. The pressure stays essentially constant and does
>not increase as the ideal gas law would predict. Similarly, what happens
>when the volume of the container is increased? More liquid vaporizes and
>the pressure remains essentially constant.
>
>One could argue that the "n" of the equation (PV = nRT) is changing since
>the number of moles of gas is changing as liquid is either condensed or
>evaporated but this is not the way the ideal gas law is commonly used. In
>this case then N(total) = n(liquid) + n(gas). The n(gas) is the "n" you
>would need to use in the ideal gas equation if you were trying to use it in
>this situation. It is usually regarded as a constant depending upon which
>particular sample of a gas you were describing. It will not be constant in
>this case.
>
>Remember the definition of an ideal gas? Among other small details, the
>individual molecules possess zero volume and they do not interact with each
>other except by colliding with each other in perfectly elastic collisions.
>There are no intermolecular attractive forces (because of this the
>nonexistent "ideal gas" cannot be liquefied). This situation is only a
>decent approximation at very low pressures (where the molecules are very
>far apart and their individual volumes are insignificant in comparison to
>the total volume of the system and interactive forces, which decrease
>markedly with distance, are insignificant) and/or very high temperatures
>(where attractive forces are very small in comparison to the average
>kinetic energies of the molecules). We are dealing with neither of these
>conditions inside this tank of N2O. Here the temperature is well below the
>critical temperature and the pressure is well above the critical pressure.
>
>(Definition {just in case you might need it}: critical temperature - the
>temperature above which a gas cannot be liquefied by pressure alone.
>Critical pressure - the minimum pressure required to liquify a gas which is
>at its critical temperature.)
>
>>Remember, the tank exploded while
>>the car was sitting in the garage...the nitrous system was not actively
>>being used (other than the bottle heater, apparently). Admittedly, there
is
>>only a small volume of gas in a full tank of nitrous oxide (mostly
liquid),
>>but even that small volume is subject to the Ideal gas law.
>
>There must be some other explanation for the explosion. A malfunctioning
>heater which caused quite high temperature? A defective or damaged tank?
>Defective or damaged plumbing, etc? A combination of more than one of
>these things? Something else?
>
>>Bob W.
>>
>>
>>
>___
> Bob
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