Xenon Di and Tetrafluoride Melting Points
Date: June 2008
Hi, I would like to ask why Xenon difluoride has a higher melting point
compared to Xenon tetrafluoride. Xenon difluoride melting point=144 C, Xenon
tetrafluoride melting point=112 C. I could not find any explanation to this as
Xenon difluoride should have a lower melting point if this is based on the fact
that Xenon difluoride has a smaller no. of electrons for polarisation and both
types of molecules are non-polar molecules. If we look from the perspective of its
molecular mass that Xenon tetrafluoride has higher molecular mass and it should has
a higher melting point, it seem to contradict the real values of the melting
points. Is there any other explanation to explain this?
Predicting melting point(s) is a tricky business. No good model for doing so exists.
Mass of the molecule plays a part, but is by no means the whole story, or even the
most important. Most metals have a high melting point compared to molecular
substances, polar molecules tend to have higher melting points than non-polar
molecules, but there are many exceptions to any of these rules.
Your rational based on an electron "count" is a single molecule argument. Using gas
phase properties/data to predict the behavior of the condensed phases
(liquid/solid) is making yet another large leap of faith.
An excellent question. The xenon fluorides have unique properties
as a group.
My own searches revealed the following:
XeF2 melts at 129°C,
XeF4 does not melt at all, but sublimes directly to gas (like CO2) at 116°C,
XeF6 melts at 49°C.
XeF4 is behaving pretty typically for a nonpolar molecule. However, the
comparison of XeF2 to XeF6 does not obey the rule of thumb that
melting points should increase with increasing molecular weight. This is usually
explained by assuming that solid XeF2 has some ionic character in the following
sense: the atoms in XeFn are believed to be highly charged. XeF2 packs itself
into a crystal structure such that each Xe "cation" is surrounded not only by its
own two Fs, but also several Fs ("anions") from neighboring molecules. This might
make the lattice energy unusually high, and therefore the melting point
would be unusually high as well. Check out the following references:
and our old friend Wikipedia, which appears to me to be correct on these three compounds
on June 12, 2008:
Best, Dr. Topper
I know XeF2 and Xe F4 sound a little like non-polar molecules such as CF4, but how
can they be?
Fluorine atoms bonds to the filled-shell, "inert gas" Xe by:
One F steals one electron outright,
then another F makes a covalent bond with the unpaired electron created by this
It is not real easy for the first F to keep its stolen electron,
but the energy of the second covalent bond helps favor the transaction.
All this tends to explain why the Fluorine counts are 2 and 4 but not 1 or 3.
Resonance, mixing of molecular electronic states, or whatever you call it,
probably equalizes the negative charge on two F's at 0.5 electron each.
Then each bond is something like 50% polar and 50% covalent.
So the Xe must be +, and the F must be -. Not really non-polar.
With 4 F's, the Xe is nearly surrounded 3-dimensionally
so the dipole moments (+..-) are less exposed to neighboring molecules.
With 2 F's, the dipoles are plainly available for broadside proximity to neighbors
of opposite orientation,
and it would be easy to make +atom to -atom proximity-links like water (H2O)
makes among itself.
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Update: June 2012