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Name: Marvin
Status: other
Grade: 9-12
Location: Singapore
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.

Vince Calder

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

Hi Marvin-

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 theft. 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. (No XeF3) 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.

Jim Swenson

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