On 31 Oct 2003 04:49:23 -0800, (basskisser) wrote:

Steven Shelikoff wrote in message ...
Rod McInnis wrote:
"basskisser" wrote in message
om...


That has absolutely nothing to do with the fact that that nitrogen
doesn't expand at the same rate as oxygen for any given temperature
change. Do you deny this?


Yes. I deny this.

Mr. Boyle denies this. Mr. Charles denies this. Mr. Gay and Mr. Lussac
deny this. They wrote laws of physics about it. Every chemistry, physics and
thermodynamics class uses these laws. Here, don't take my word for it, let's
take a look at some of the information available from the net.

As an example:

Department of Chemistry
California State University, Sacramento
http://kekule.chem.csus.edu/gaslaws

You might as well give up now. No matter how much proof you provide,
basskisser will not believe it and will find some way to weasel out,
probably by saying he answered you in some other post without being able to
show where.

The thing I find amazing is that he claims to be a structural engineer, or
something like that. A professional engineer even. He is a sad testament
to whatever college he graduated from. This is basic high school physics
that is also reempasized in first year college engineering physics. For
him to be so blatently wrong in something that is so provable that he's
wrong and still not be able to admit it is just plain sad.

Boyles' law states that for a GIVEN GAS, the rate of expansion versus
temperature is inversely proportional. FOR A GIVEN GAS. It does not
state, however, that one gas expands as temp. increases, at the same
rate as another gas.

Very good. Now look up Charles' Law, V1/T1=V2/T2. Then the ideal gas
law, PV=nRT. And note that at low pressures (i.e., the pressures at
which are normally inside tires) all gasses act like ideal gasses and
have the SAME pressure vs. temperature vs. volume relationship..

Your phase change argument elsewhere is complete crap if it's meant to
show that nitrogen has a different PVT relationship vs. oxygen vs. air
at the pressures we're discussing. Now, I'll be nice and give you a
hint on how you could possibly defend your statement about air having a
less linear pressure temperature relationship then nitrogen: Look up
how much more or less ideal air is vs. nitrogen.

But even that argument has a problem in that 1) both of them behave like
an ideal gas at the low pressures we're talking about, where the
molecules are still far enough apart that their attraction doesn't
change the linear relationship PV=nRT. And 2) air is mostly nitrogen
anyway. So if you put both of them at a pressure where they deviate
from the ideal gas law (a very high pressure that would blow up most
storage tanks let alone any tire) they would both deviate from the ideal
gas law almost equally.

Keep trying.

Steve