On Tue, 22 Jul 2003, Doug Jones wrote:
> As I vaguely recall, it has something to do with the attractive Van der
> Walls forces- which are what make gases non-ideal. Expansion through an
> orifice, even when it does no macroscopic work, does work against the
> van der walls force.
That sounds right, and a quick look at my old Phys Chem text reveals a
discussion of "internal pressure" due to intermolecular forces within the
gas -- first measured via the Joule-Thomson effect -- which matches up
with that.
Intermolecular forces are one of the two things that make gases non-ideal,
actually. The other is finitely small molecular volume, which becomes
significant as pressure rises -- only the volume *not* occupied by the
molecules themselves is available for PV=nRT behavior. Molecules of an
ideal gas would interact only with the container walls, not with each
other, and would have zero volume. The first step up in realism from the
ideal-gas law is the van der Waals law:
(P + an^2/V^2) * (V - bn) = nRT
where a and b are gas-specific empirically-derived constants adjusting for
intermolecular forces and molecular volume respectively. b is pretty
obvious, it's just the molecular volume per mole; a is a more empirical
fudge factor. You can see that, indeed, the intermolecular-forces term
adds to the pressure (assuming a is positive) -- the gas acts as if it's
under greater pressure, thus is more compact than it should be.
(There are alternatives, e.g. the Berthelot equation in which the
fudge-factor's denominator is TV^2, which may do better, depending on
the conditions.)
It's not surprising that hydrogen and helium are the odd men out for J-T
inversion temperature; as I noted before, they're exceptions to a lot of
rules. They -- and sometimes neon, which is a borderline case -- are
called the "quantum gases", because quantum physics is quite significant
in their behavior.
Henry Spencer
[EMAIL PROTECTED]
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