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Boiling Points of Gasses
Name: Tina A.
Status: educator
Age: 30s
Location: N/A
Country: N/A
Date: Sunday, December 08, 2002
Question:
I would like to know why oxygen gas has a higher boiling
point than nitrogen and fluorine gases.
Replies:
First, it is not easy to pinpoint a "cause" of why certain substances have
a certain boiling point (or melting point). Evaporation is a complicated
process.
Second, the boiling point is not the correct measure to use if one is
comparing volatility. The "normal" boiling point is the temperature at which
the vapor pressure is 1 atm (760mm of Hg).
This selection of the temperature is arbitrary. You might choose to compare
the temperature where the vapor pressure is 100 mm of Hg, or 1000mm of Hg.
If you would do so you would find that some of the ordering of these
"boiling points" change. So boiling point is not a good characteristic
parameter of the volatility.
Third, consider the process that is occurring. molecules of liquid have
sufficient volatility that a certain number of them can have sufficient
kinetic energy to escape the liquid phase -- of course the reverse is also
happening and some vapor molecules are striking the surface of the liquid
and "sticking". It should be evident that the higher the number of molecules
that have this energy the more this process is occurring -- so evaporation
has something to do with the vapor density.
This would suggest that the critical temperature (the temperature at which
the density of the liquid and vapor are equal) would be a better parameter
to compare.The critical temperature of N2, O2, and F2 is 126.2, 154.6, and
144.3 (all temperatures in K). Molecular size, shape, volume, mass, and
intermolecular forces all play a role in determining the volatility of a
substance. The average is 141.7 K, so the critical temperature of all three
gases is the same within 10%, and that's about all the sensitivity you can
expect.
The critical temperature of the elements Ne, Ar, Kr, and Xe is 44.4, 150.8,
209.4 and 289.7. These follow a mass trend, but that may be coincidental,
because the atomic size of these molecules (all with the same spherical
shape) should also play a role. The atoms have very different polarizability
(you can think of this as how "squishy" the atoms are), and this plays an
important role in the formation of induced dipoles, which plays a role in
the volatility.
The bottom line is: It is OK to make these property comparisons to help
you develop an intuition about how atoms and molecules behave, but do not
take them too seriously, and expect some deviations. Most all the property
comparisons you make depend upon the interplay of several parameters, so it
is not possible (usually) to say thus and so is the "cause" of the trend.
Nonetheless, chemists do it all the time and come up with the right answer
or trend for the wrong reason!!!
Vince Calder
Dear Tina,
Rules are meant to be broken!!
F2 has more electrons than O2, which has more electrons than N2.
If more electrons ALWAYS meant greater dispersion forces then
the boiling point ordering would be N2 < O2 < F2. But the REAL
property which makes dispersion forces large or small is POLARIZABILITY
(this means how easy or hard it for electrons to be pushed around
within the molecule by outside forces).
Since polarizability often increases with the number of electrons
you might guess the above ordering to be correct. However, something
else is also changing; the effective nuclear charge is increasing
as you go from N2 to O2 to F2, which tends to DECREASE the
polarizability. O2 probably has the optimum combination; it has more
electrons than N2, but it is more polarizable than F2 (F is the most
electronegative element, after all).
It is generally safer to make comparisons within the same group
(like comparing H2S to H2Se to H2Te) or within a series of otherwise
similar compounds (like methane, ethane, propane, butane, pentane...).
Then the increasing number of electrons correlates nicely with the
polariability, which almost always correlates with the boiling point
(but please notice that I said "almost").
prof. topper
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Update: June 2012
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