| Pages:
1
2 |
ave369
Eastern European Lady of Mad Science
Posts: 596
Registered: 8-7-2015
Location: No Location
Member Is Offline
Mood: No Mood
|
|
A probably very stupid idea
A probably very stupid idea just hit me, but it may seem interesting. What happens if I burn a PTFE flash powder in a glovebox filled with xenon?
As far as I heard, the extremely high temperatures produced by the flash powder cause fluorine to liberate from PTFE, which then immediately reacts
with the metal in the flash powder. So, if xenon is also present, will any xenon fluorides form? If it works, it may be a relatively safe and
available way to produce xenon compounds!
Maybe a minute amount of xenon tetrafluoride forms, sublimes and crystallizes on the glovebox walls. So if I burn a lot of flash powder in the box and
then wash it with a solution of sodium hydroxide (to immediately bind both xenon trioxide and hydrogen fluoride which form when xenon tetrafluoride
reacts with water), the resulting solution will have sodium xenate?
[Edited on 5-8-2015 by ave369]
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
I do not see any practical benefits from this. I certainly can imagine that a tiny fraction of fluorine atoms bonds with Xe atoms, but in the horrible
mess of all kinds of compounds you would have a hard time detecting any Xe-compounds, let alone isolating a macroscopic useful amount of Xe-compounds.
I myself have looked into obtaining Xe-compounds, but they are very very expensive. There is one seller in NL who is willing to sell XeF2 to amateur
chemists, but the price is appr. EUR 185 per gram, appr. EUR 700 for 5 grams. The only other commercially available Xe-compound is Na4XeO6 (sodium
perxenate), but this is not in the seller's list of available chemicals. Xenon itself also is very expensive, a glovebox full of Xe at atmospheric
pressure probably will cost hundreds of euros.
|
|
|
ave369
Eastern European Lady of Mad Science
Posts: 596
Registered: 8-7-2015
Location: No Location
Member Is Offline
Mood: No Mood
|
|
In my country, cylinders of medical xenon can be obtained on the gray market cheaply: a small cylinder of medical grade xenon-oxygen mix, with enough
xenon to fill one small glovebox, can be bought for 2000 rubles (this is something like $30-$40). It is defitsit, which means it is hard to
find, but it isn't prohibitely expensive.
Also, there may be a bajillion of fluorine compounds forming, but only a few of them are water-soluble. The lighter (the most toxic) perfluorocarbons
will be ejected from the glovebox when you turn on the ventilation, together with unreacted xenon. The heavier, non-volatile perfluorocarbons are
insoluble or immiscible with water, and will not react with the alkali; it would be easy to separate them. So is carbon soot. The washing solution
will contain only several substances: the original NaOH, the fluoride of the metal used in the flash powder (if it's magnesium, the fluoride will be
almost insoluble), and the products of reaction between xenon tetrafluoride and the alkali: sodium fluoride and sodium xenate. That's only four
chemicals, and it will be possible to detect xenate in this solution, I think.
[Edited on 5-8-2015 by ave369]
|
|
|
careysub
International Hazard
Posts: 1339
Registered: 4-8-2014
Location: Coastal Sage Scrub Biome
Member Is Offline
Mood: Lowest quantum state
|
|
Why not do this in a reaction vessel (aka a "pipe") instead of burning it the "open xenon environment" (can't say "open air") in a glove box?
Pressurizing it might enhance the effect also.
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
| Quote: | | As far as I heard, the extremely high temperatures produced by the flash powder cause fluorine to liberate from PTFE |
I would like to know more about this. Do you have a reference, please?
Polymers can thermally crack to form their monomers, so one could consider the hypothetical case where PTFE thermally cracks to form
tetrafluoroethylene (TFE). Not saying this is the only cracking product, I just want to consider the thermodynamics of a likely product.
Certainly a reaction with aluminium can generate the high temperature required for thermal cracking, so I don't think there's much one needs to
consider further there.
Let us then consider the 'combustion' of the idealised cracking product TFE in a Xe/O2 stream, assuming you run your thermite aluminium-lean in order
to produce it.
C2F4(g) + 2O2(g) + 2Xe(g) => 2CO2(g) + 2XeF2(s)
The dG_rxn_standard = -339kJ/mol. C2F4
dG << 0
So that checks the first thermodynamics box.
References:
(1) XeF2 thermodynamic properties obtained in the abstract from http://scitation.aip.org/content/aip/journal/jcp/57/11/10.10...
(2) Other thermodynamic properties obtained from the NIST webbook.
|
|
|
Nicodem
Super Moderator
Posts: 4230
Registered: 28-12-2004
Member Is Offline
Mood: No Mood
|
|
And what about the expected reaction:
C2F4(g) + O2(g) => 2COF2(g)
Why would xenon just magically take over fluorine atoms? This does not make any redox sense.
|
|
|
aga
Forum Drunkard
Posts: 7030
Registered: 25-3-2014
Member Is Offline
|
|
Why do you want Xenon compounds ?
There may be better/easier routes to a Xenon compound, or an analogue you use instead (if you have a specific reaction in mind).
[Edited on 5-8-2015 by aga]
|
|
|
ave369
Eastern European Lady of Mad Science
Posts: 596
Registered: 8-7-2015
Location: No Location
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by deltaH  | | Quote: | | As far as I heard, the extremely high temperatures produced by the flash powder cause fluorine to liberate from PTFE |
I would like to know more about this. Do you have a reference, please? |
Google "cf4 thermal decomposition" and you will see bajillions of references. CF4 decomposes and liberates fluorine at 2300-2800 C (2500-3000 K).
Fluorine flash powders do provide such an extreme temperature.
| Quote: | | Why do you want Xenon compounds ? |
One word: perferrate. I've read some Russian papers which argue that ferrate (VIII) exists and want to give it a try. But to synthesize it, I need
bigger oxidizers than ferrate (VI). Xenates and perxenates seem to pack enough punch. Maybe perbromate will do the trick, but to synthesize it you
need perxenate.
[Edited on 5-8-2015 by ave369]
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
| Quote: | And what about the expected reaction:
C2F4(g) + O2(g) => 2COF2(g)
Why would xenon just magically take over fluorine atoms? This does not make any redox sense. |
It's an assumption that CO2 forms instead of COF2 at elevated temperature. As I said, it checks the first thermodynamics box or should I say
the first easy to calculate back-of-the-envelope box  What one now needs to look
at is what happens to the thermodynamics as a function of temperature and validate the assumption.
There isn't Cp data as a function of temperature available for XeF2 at high T AFAIK, but what you raise can essentially be simplified to the
thermodynamic question around the following hypothetical equilibrium for which there is sufficient data on the NIST webbook site:
2COF2 + O2 <=> 2CO2 + 2F2
Since the dGf of the elements is zero and the same stoichiometric factor applies for COF2 and CO2, you can compare dGfs of those two species as a
function of temperature and see at what temperature dGf_COF2 > dGf_CO2, if any.
That said, experience tells that the RHS wins at high temperature, although it might be interesting to know what the minimum temperature is.
[Edited on 5-8-2015 by deltaH]
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
| Quote: Originally posted by ave369 | | Quote: Originally posted by deltaH | | Quote: | | As far as I heard, the extremely high temperatures produced by the flash powder cause fluorine to liberate from PTFE |
I would like to know more about this. Do you have a reference, please? |
Google "cf4 thermal decomposition" and you will see bajillions of references. CF4 decomposes and liberates fluorine at 2300-2800 C (2500-3000 K).
Fluorine flash powders do provide such an extreme temperature. |
Thanks, very interesting reading!
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by deltaH |
Ah, as I said, it checks the first thermodynamics box or should I say the first easy to calculate back-of-the-envelope box
What one now needs to look at is what happens to the thermodynamics as a function of temperature.
There isn't Cp data as a function of temperature available for XeF2 at high T AFAIK, but what you raise can essentially be simplified to the
thermodynamic question around the following hypothetical equilibrium for which there is sufficient data on the NIST webbook site:
2COF2 + O2 <=> 2CO2 + 2F2
Since the dGf of the elements is zero and the same stoichiometric factor applies for COF2 and CO2, you can compare dGfs of those two species as a
function of temperature and see at what temperature dGf_COF2 > dGf_CO2, if any.
[Edited on 5-8-2015 by deltaH] |
No. That's incorrect.
ΔG<sub>reaction</sub><sup>0</sup> (at STP) can becompletely evaluated at 298 K, even if (obviously!) during an
exothermic reaction temperature evolves. As long as as the system ends back up at 298 K, the functions ΔG<sub>f</sub>(T) of the species
don’t matter because the equilibria adjust back to 298 K. Of course none of this says anything about kinetics at all.
But I’m not sure why I’m contributing here: this mad cap scheme to ‘prepare’ Xe fluorides as suggested by the OP will never be evaluated and
is as such YET another SM red herring.
And this bit:
| Quote: Originally posted by ave369 |
As far as I heard, the extremely high temperatures produced by the flash powder cause fluorine to liberate from PTFE, which then immediately reacts
with the metal in the flash powder. So, if xenon is also present, will any xenon fluorides form? If it works, it may be a relatively safe and
available way to produce xenon compounds!
|
...shows such a profound lack of understanding of equilibria it's hard to think where to begin to debunk this skunk. In plain English, it's nonsense.
[Edited on 5-8-2015 by blogfast25]
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
| Quote: Originally posted by blogfast25 | | Quote: Originally posted by deltaH |
Ah, as I said, it checks the first thermodynamics box or should I say the first easy to calculate back-of-the-envelope box
What one now needs to look at is what happens to the thermodynamics as a function of temperature.
There isn't Cp data as a function of temperature available for XeF2 at high T AFAIK, but what you raise can essentially be simplified to the
thermodynamic question around the following hypothetical equilibrium for which there is sufficient data on the NIST webbook site:
2COF2 + O2 <=> 2CO2 + 2F2
Since the dGf of the elements is zero and the same stoichiometric factor applies for COF2 and CO2, you can compare dGfs of those two species as a
function of temperature and see at what temperature dGf_COF2 > dGf_CO2, if any.
[Edited on 5-8-2015 by deltaH] |
No. That's incorrect.
ΔG<sub>reaction</sub><sup>0</sup> (at STP) can becompletely evaluated at 298 K, even if (obviously!) during an
exothermic reaction temperature evolves. As long as as the system ends back up at 298 K, the functions ΔG<sub>f</sub>(T) of the species
don’t matter because the equilibria adjust back to 298 K. Of course none of this says anything about kinetics at all.
But I’m not sure why I’m contributing here: this mad cap scheme to ‘prepare’ Xe fluorides as suggested by the OP will never be evaluated and
is as such YET another SM red herring.
[Edited on 5-8-2015 by blogfast25] |
The issue nicodem raised is that COF2 will be forming and not CO2. I am saying that at low temperature COF2 will indeed form (and is
the thermodynamically preferred product), but that at high temperature CO2 is the preferred thermodynamic product and gave an illustrative example of
the equilibrium at play.
Hence this is why I made that assumption initially in calculating the thermodynamics at standard conditions that the reaction was forming CO2, because
considering the reaction to form COF2 would be meaningless as there is no release of the fluorides and reaction to consider (the point
nicodem I believe was trying to make, but my point is that that is not the case at high temperature).
I've edited the original to emphasise that it's an assumption of what's going on at high temperature, hopefully that's clearer.
*************
Ah wait, I think I understand what you mean (maybe) is that CO2 and F2 (XeF2) will form COF2 when you cool it later on? If so then yes, that will be
favoured, which is why you would need to quench the reaction as fast as possible and hope the kinetics for the reverse reaction is slow enough at low
temperature to salvage product.
Anyway, I'm thoroughly confused and my brains addled, so going to sleep now...
I'm even writing CO2F2 
You and nicodem may well be right, fluorine kinetics are after all usually very fast, so perhaps practically one cannot quench the
reaction fast enough to isolate the high temperature product and will simply end up with the low temperature thermodynamically favoured COF2.
[Edited on 5-8-2015 by deltaH]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
And where, pray tell, would the fluorine come from, huh? The whole idea of the teflon flash powder is that PTFE reacts with the metal, forming metal
fluoride. I.o.w the fluorine simply gets gobbled up, fuelling the enthalpy released by the flash powder. So the only potential source of fluorine
would be extra teflon. Who can say whether that will decompose to C/F<sub>2</sub> though? Why can't it just melt, depolymerise somewhat
and volatilise?
Seriously, this whole scheme is like hoping to use oxygen from a thermite reaction!
Keep doodling but you should know better, deltaH...
[Edited on 5-8-2015 by blogfast25]
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
Yes, my gut feeling was also against this, but I have to doodle 
My mistake was making the assumption that the reaction forms CO2. Even if the assumption is justified at high temperature, it may well not quench fast
enough and reverse while cooling, but that's a kinetics issue.
I don't know if your question about where the fluorine would come from was directed at me or ave369, but I interpreted this as being
exactly as you said, one uses substoichiometric amounts of Al and cracks the excess PTFE into volatile fluorocarbons that then 'burn' in a oxy-xenon
flame.
The aluminium is sitting in molten PTFE, so it doesn't see the oxygen, it just sees PTFE which is will happily react with. This releases cracked
fluorocarbons.
The question is can the cracked fluorocarbons be 'burnt' in an oxy-xenon flame and quenched fast enough to form CO2 and XeF2. I simply considered the
very first basic step on the path to that, namely, the thermodynamics of that reaction, specifically the high-temperature version.
Again, whether it can survive the cool down without forming COF2 is another matter altogether.
[Edited on 5-8-2015 by deltaH]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by ave369 |
As far as I heard, the extremely high temperatures produced by the flash powder cause fluorine to liberate from PTFE, which then immediately reacts
with the metal in the flash powder. So, if xenon is also present, will any xenon fluorides form? If it works, it may be a relatively safe and
available way to produce xenon compounds!
|
The OP's idea probably comes from a misinterpretation of a representation of Hess' Law to calculate ΔG:
PTFE == > carbon + F2
Me + F2 === > MeF2
PTFE + Me === > carbon + MeF2
But this is only an 'imaginary' pathway to allow calculation of ΔG, it has nothing to do with how this reaction actually proceeds. No free
F<sub>2</sub> is ever present in the reaction mix.
|
|
|
ave369
Eastern European Lady of Mad Science
Posts: 596
Registered: 8-7-2015
Location: No Location
Member Is Offline
Mood: No Mood
|
|
Okay, let's do it differently. Let's not add xenon to the original flash powder, let's put some excess PTFE and surround the powder with iron instead.
A mix of aluminium and more PTFE than stoichiometry with aluminium requires, and an outer layer of powdered iron. Iron becomes fluorinated together
with aluminium and forms FeF3.
Later we make another powder, with FeF3 instead of PTFE. This compound definitely does decompose and liberate fluorine in macroscopic amounts at high
temp, and requires lower temp to do this. And it's this we put into a pressurized bomb with xenon.
Will this work?
Of course, all powders are lit remotely, from 100 meters in an uninhabited locale. It's fluorine we are dealing with, after all.
[Edited on 6-8-2015 by ave369]
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
| Quote: Originally posted by ave369 | Okay, let's do it differently. Let's not add xenon to the original flash powder, let's put some excess PTFE and surround the powder with iron instead.
A mix of aluminium and more PTFE than stoichiometry with aluminium requires, and an outer layer of powdered iron. Iron becomes fluorinated together
with aluminium and forms FeF3.
Later we make another powder, with FeF3 instead of PTFE. This compound definitely does decompose and liberate fluorine in macroscopic amounts at high
temp, and requires lower temp to do this. And it's this we put into a pressurized bomb with xenon.
Will this work?
Of course, all powders are lit remotely, from 100 meters in an uninhabited locale. It's fluorine we are dealing with, after all.
[Edited on 6-8-2015 by ave369] |
No this is less likely to work. The problem is the very strong C-F bond. You need either a very strong oxidant (e.g. oxygen) or a very strong
reductant (e.g. aluminium) to persuade carbon to part with its fluorine.
The formation of CO2 and very high temperature is a powerful thermodynamic driver and so a special case. That's why the oxy-xenon atmosphere had some
small hope.
However, FeF3 is thermodynamically highly unfavoured at high T! The reverse reaction would likely readily occur, i.e. you could fluorinate carbon with
FeF3 to generate fluorocarbons and FeF2/Fe!
[Edited on 6-8-2015 by deltaH]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by ave369 |
Later we make another powder, with FeF3 instead of PTFE. This compound definitely does decompose and liberate fluorine in macroscopic amounts at high
temp, and requires lower temp to do this. And it's this we put into a pressurized bomb with xenon.
|
Really? The Standard Heat of Formation of FeF<sub>3</sub> is ΔH<sub>f</sub><sup>0</sup> = - 1042 kJ/mol (NIST
webbook). That requires very high temperature to get significant dissociation. At that temperature any [Xe,F] compound would be completely
dissociated. On cooling the Fe and F would of course simply recombine. THAT's what the equilibria tell us.
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
Oops, I assumed FeF3 was like CoF3 by what ave369 was describing, my bad.
| Quote: | | However, FeF3 is thermodynamically highly unfavoured at high T! |
Then this is junk, sorry 
|
|
|
ave369
Eastern European Lady of Mad Science
Posts: 596
Registered: 8-7-2015
Location: No Location
Member Is Offline
Mood: No Mood
|
|
Okay. Then I'm abandoning the project and not going to do anything involving fluorine at all. As I thought, the idea is stupid.
[Edited on 6-8-2015 by ave369]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by deltaH | Oops, I assumed FeF3 was like CoF3 by what ave369 was describing, my bad.
| Quote: | | However, FeF3 is thermodynamically highly unfavoured at high T! |
Then this is junk, sorry
|
Dunno about CoF3 but for CoF2 it's largely junk too: HoF = - 672 kJ/mol, still quite respectable.
Really this whole 'idea' is predicated on the idea that PTFE + Xe ===> C + [Xe,F] over PTFE + metal === > C + metal fluoride and that's absurd,
considering the low thermal stability of Xe fluorides (compared to metal fluorides).
[Edited on 6-8-2015 by blogfast25]
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
As for the CoF3 forming fluorine on heating... it's something woelen once mentioned in passing on the reagent purchases thread,
perhaps he didn't mean elemental fluorine, but fluorination? Anyway, wiki says it's a strong fluorinating agent. I just assumed that FeF3 behaved
similarly because of what ave369 said, never bothered to check, my bad, sorry.
|
|
|
Dan Vizine
National Hazard
Posts: 628
Registered: 4-4-2014
Location: Tonawanda, New York
Member Is Offline
Mood: High Resistance
|
|
What's wrong with the accepted methods to Xenon fluorides? They actually work.
"All Your Children Are Poor Unfortunate Victims of Lies You Believe, a Plague Upon Your Ignorance that Keeps the Youth from the Truth They
Deserve"...F. Zappa
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Nothing. But this thread isn't about that: it's about someone's bandwidth consuming 'idea'! 
|
|
|
halogen
Hazard to Others
Posts: 372
Registered: 18-4-2004
Member Is Offline
Mood: No Mood
|
|
http://onlinelibrary.wiley.com/doi/10.1002/1521-3757%2820011...
Durch Fluor-Chlor-Austausch an XeF+ entsteht das unterhalb von −10°C stabile, orangefarbene Chlor-Xenon-Salz XeCl+Sb2F11− (siehe Struktur im
Kristall). Dies ist erst das dritte Edelgas-Monohalogen-Kation, und wie Rechnungen zeigen, dürfte die Synthese weiterer derartiger Kationen – mit
Ausnahme der ArF+-Spezies – auch nicht gelingen.
Xenon Dichloride is too unstable to be isolated. It may be, however, that argentous fluoride and dichlorine prove possible to oxidize Xe. I don't feel
like looking up tables, just "throwing the idea out there". If it helps. The original preparation of IF5 by George Gore - prior to Moissan's
preparation of fluorine, and what a badass Gore was by the way, was from AgF and I2 under heavy heating.
F. de Lalande and M. Prud'homme showed that a mixture of boric oxide and sodium chloride is decomposed in a stream of dry air or oxygen at a red heat
with the evolution of chlorine.
|
|
|
| Pages:
1
2 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
