deltaH
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tetrafluoroammonium perchlorate superoxidant
We do so love speculating a little on super oxidants on SM don't we? I'll keep it short and sweet and propose:
tetrafluoroammonium perchlorate, NF4ClO4
... and the proposed synthetic route?
More or less standard fluorine electrolysis cell using molten KF.2HF at 70C, but with a Nafion membrane to separate anolyte from the catholyte and
adding ammonium perchlorate at the anode side.
That way as the fluorine forms, it almost immediately reacts with the ammonium ions to form progressively more fluorinated amines up until the
tetrafluoroammonium species.
Since NF4+ and ClO4- are similarly sized, both tetrahedral and non-polar, this salt would probably have excellent crystal packing, high density and
low solubility, thus precipitating from the KF.2HF mix.
I would say this would be the fluoro analogue of the much better known nitronium perchlorate as trifluoramine has been described as being similar in
chemical reactivity and toxicity to nitrogen dioxide, but still a strong oxidant and slow fluorinating agent.
Unlike NO2+ though, I would suspect that NF4+ is significantly less reactive at room temperature because the nitrogen atom is coordinatively saturated
by the poorly labile fluoride 'ligand'.
[Edited on 30-10-2013 by deltaH]
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vulture
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| Quote: |
We do so love speculating a little on super oxidants on SM don't we? I'll keep it short and sweet and propose:
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Referenced speculation would be nicer.
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Unlike NO2+ though, I would suspect that NF4+ is significantly less reactive at room temperature because the nitrogen atom is coordinatively saturated
by the poorly labile fluoride 'ligand'.
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You can not "coordinatively saturate" nitrogen with fluorine, because NF4- is not a complex but a sigma bonded species.
Faulty nomenclaturic diarrhea does not make your ideas more credible, on the contrary.
One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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DubaiAmateurRocketry
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Yes the it is extremely energetic, but flourine is too exothermic.
Also, flourine does not actually burn that well with hydrogen and aluminum producing almost same heat as oxygen, in that case oxygen is prefered since
H2O expands more. Flourine is burns at the same temperature as aluminum than aluminum does with oxygen.
Flourine burns slightly hotter with magnesium than aluminum, which is much more expensive and less dense than aluminum.
For hottest flourine reactions, it would be Flourine and lithium, beryllium probably.
According to wikipedia and some other sites, a rocket engine using F + molten Li reached a specific impulse of 500. but its extremely unstable, and
they failed so many times.
Also flourine would be toxic.
If your going for superoxidants,
then i think Nitronium - Dinitramide would be good, aka - trinitramide.
I wonder if it is possible to have a tetra-nitramide cation.
[Edited on 30-10-2013 by DubaiAmateurRocketry]
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deltaH
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@vulture
Requested paper and will append. Your ion isn't correct, it's supposed to be NF4+ not minus.
@DubaiAmateurRocketry
This is not about fluorine liquid or gas, but a solid salt.
NF4ClO4 has an extremely high oxidant mass content, if you sum fluorine and oxygen atoms, it comes to 74 wt. %. That's impressive, especially if the
compound is stable, which it might well be.
If it turns out that it can be prepared by the electrochemical fluorination of ammonium perchlorate, then it is practically accessible at not too high
a production cost.
In combustion with polymeric organic fuels + Al, it would probably produce AlF3, N2, HCl, CO2, CO and H2O if you have sufficient Al present. If so,
AlF3 is harmless.
One of the crappy things of NH4ClO4 is the ammonium ion... ammonia isn't a good fuel compared to hydrocarbons, so losing 1.5 of your oxygens to form
H2O here actually detracts from the total energy you can release in combination with a fuel. In NF4ClO4, you ditch the useless ammonia bit and replace
it with another energetic oxidant...a fluorinated nitrogen trifluoride! This allows you to get much more energy in combination with fuels because you
don't have to 'waste' some with the oxidation of ammonia.
| Quote: | then i think Nitronium - Dinitramide would be good, aka - trinitramide.
I wonder if it is possible to have a tetra-nitramide cation. |
It is highly unlikely that nitronium dinitramide could be isolated, let alone be stable. Both the dinitramide and nitronium ions are way to reactive,
they would likely self react rapidly to form NOx and N2. Rapidly react here is more a euphemism for explode 
While I have even greater doubts about the tetranitramide cation, I think you should request it be molecularly modelled in the computational chemistry
section, should give you some initial idea if this has any hope of being stable.
[Edited on 30-10-2013 by deltaH]
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DubaiAmateurRocketry
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| Quote: |
It is highly unlikely that nitronium dinitramide could be isolated, let alone be stable. Both the dinitramide and nitronium ions are way to reactive,
they would likely self react rapidly to form NOx and N2. Rapidly react here is more a euphemism for explode
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But it is stable.
http://en.wikipedia.org/wiki/Trinitramide
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deltaH
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I see, but I'd be wary to link "was detected" to is stable, maybe we should request that paper. Seems like it might simply have been detected in small
amounts in a strongly nitrating solution (conditions under which dinitramine might be prepared).
If they isolated it in pure form and did some stability experiments, then I will be very excited that a revolutionary new oxidant is on it's way!
My gut feeling tells me this will be too reactive and/or unstable in pure form to be a practical oxidant, but I may be wrong. Try to get a hold of
that paper and read it critically.
Better yet, post it so that we can all read up about, but lets' keep it detached from this thread, I want to consider tetrafluoroammonium perchlorate
in this thread.
I would suggest you start a thread about it and reference that paper.
Thanks.
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DubaiAmateurRocketry
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Quote: Originally posted by deltaH  |
If they isolated it in pure form and did some stability experiments, then I will be very excited that a revolutionary new oxidant is on it's way!
My gut feeling tells me this will be too reactive and/or unstable in pure form to be a practical oxidant, but I may be wrong. Try to get a hold of
that paper and read it critically.
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Well i remember i read a paper about it, but not in depth.
It was isolated and is stable.
However the sad thing is.. It is a liquid. and becomes gas readily at RT.
http://thehuwaldtfamily.org/jtrl/research/Propulsion/Rocket%...
According to them, it is 19 percent denser than AP which makes it 2.32g/cm2
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DubaiAmateurRocketry
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And as for the tetrflouroammonium cation, i would guess the nitrate or dinitramide would be better ? since there are already enough oxidants and
chlorine is not needed.
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woelen
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NF4ClO4 is stable and was synthesized somewhere around 1980 by K.O. Christe and W.W. Wilson. This Christe-person is quite well known. He has done a
lot of extremely esoteric chemistry and even tried to make the salt N10 (pentazolium pentazolide), but this did not succeed.
The synthesis of NF4ClO4, however, was very difficult and involved reacting NF3, F2 and several other chemicals. There is no known cheap and easy
route to NF4ClO4, so it will not be more than an insanely expensive lab-curiousity. Nitrogen has oxidation state +5 in the NF4(+) ion and this can
only be achieved by oxidizing NF3 with F2, which yields NF4F (tetrafluoroammonium fluoride) as intermediate, which can be reacted with other chemicals
to obtain salts of the NF4(+) ion.
The ion NF4(+) is extremely reactive and in no way can be compared with NH4(+).
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deltaH
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| Quote: Originally posted by DubaiAmateurRocketry | | Quote: Originally posted by deltaH |
If they isolated it in pure form and did some stability experiments, then I will be very excited that a revolutionary new oxidant is on it's way!
My gut feeling tells me this will be too reactive and/or unstable in pure form to be a practical oxidant, but I may be wrong. Try to get a hold of
that paper and read it critically.
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Well i remember i read a paper about it, but not in depth.
It was isolated and is stable.
However the sad thing is.. It is a liquid. and becomes gas readily at RT.
http://thehuwaldtfamily.org/jtrl/research/Propulsion/Rocket%...
According to them, it is 19 percent denser than AP which makes it 2.32g/cm2
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Ah so similar sort of problem to tetranitromethane then... makes sense, but I am still very pleasantly
surprised to learn that it's stable.
The increase in density is of course not surprising. I would expect similar high density from NF4ClO4.
Incidentally, if your tetranitroammmonium cation is plausible, then you might be able to synthesis it from two known compounds (one of which you
stated exists and is stable):
N(NO2)3 + N2O5 => N(NO2)4+NO3-
I'd still favour NF4ClO4 though if and only if it can be synthesised by electrochemical fluorination of ammonium perchlorate, simply for practicality
sake and stability. Such a route can be mass produced, N(NO2)4+NO3- is highly unlikely to be so... as pretty as it may be
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deltaH
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| Quote: Originally posted by woelen | NF4ClO4 is stable and was synthesized somewhere around 1980 by K.O. Christe and W.W. Wilson. This Christe-person is quite well known. He has done a
lot of extremely esoteric chemistry and even tried to make the salt N10 (pentazolium pentazolide), but this did not succeed.
The synthesis of NF4ClO4, however, was very difficult and involved reacting NF3, F2 and several other chemicals. There is no known cheap and easy
route to NF4ClO4, so it will not be more than an insanely expensive lab-curiousity. Nitrogen has oxidation state +5 in the NF4(+) ion and this can
only be achieved by oxidizing NF3 with F2, which yields NF4F (tetrafluoroammonium fluoride) as intermediate, which can be reacted with other chemicals
to obtain salts of the NF4(+) ion.
The ion NF4(+) is extremely reactive and in no way can be compared with NH4(+). |
I remember reading about this N10 story a long time ago. Good luck to him... maybe he's just dreaming to give nitrogen a smell
Some questions which I will look into:
Has the electrochemical fluorination of ammonia been researched?
If not then an attempt could be made to electrolyse NH4HF2 (m.p 126C), which is readily commercially available as an etchant.
If this evolves NF3(g) at the anode, as I expect it could, then there is a strong case to be made towards preparing NF4ClO4 by the electrochemical
fluorination of NH4ClO4 in molten fluoride media.
[Edited on 1-11-2013 by deltaH]
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Dany
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There is two opinion for naming NF4+ cation. First, the synthesis of NF4+ has been reported by two groups in the same year
(1966) and in the same journal (Inorganic and Nuclear Chemistry Letters) see (Inorganic and Nuclear Chemistry Letters Volume 2, Issue 3,
March 1966, Pages 83–86 and Inorganic and Nuclear Chemistry Letters Volume 2, Issue 3, March 1966, Pages 79–82) the team of K.O. Christe says that
the name tetrafluoroammonium is inappropriate:
"We have named the NF4+ cation tetrafluoronitronium(V), as a derivative of hypothetical NF5. The name tetrafluoroammonium is
inappropriate due to the polarity of N-F bond"
however, the team of M.E. Hill called it simply tetrafluoroammonium.
Dany.
[Edited on 1-11-2013 by Dany]
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deltaH
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Thanks Dany for that insight into this compound's nomenclature. Personally, I agree with Christe that "tetrafluoronitronium(V)" is
much more appropriate.
I will refer to this compound henceforth as tetrafluoronitronium(V) perchlorate or abbreviate it as TFNP.
*****
I've just realised the first big problem with synthesising TFNP by employing a hypothetical electrochemical fluorination route.
That is that you would form NF3(g) at the anode instead of further oxidising to NF4+!
NF3 has a b.p. of -129C and practically non-polar (dipole moment is 0.234D), so I would expect it only to dissolve in molten hot salts in the couple
of ppm level.
Electrolysis of ammonium salts in anhydrous fluoride media would thus most likely form NF3(g).
Now, while it is known that you can react NF3(g) with F2 to form tetrafluoronitronium fluoride, the problem is that when you carry out the
electrolysis, if you have ammonium ions in solution (which you would), you have no chance of forming F2 and so probably no chance of forming NF4+.
You need to carry out a second electrochemical fluorination using non ammonium containing salts, for example the classic KF.2HF melt, together with a
porous carbon anode through which is diffused NF3(g) from the centre. As such a setup is expected to form fluorine gas, you have a chance of forming
NF4+ in solution as your anode product (in fact more than a chance).
But this can only yield tetrafluoronitronium(V) fluoride and not the perchlorate. This is another big problem!
****
Another point is that according to wiki, Christe prepares fluorine perchlorate gas by thermal decomposition of NF4ClO4, the question is what is the
thermodynamics of this like... specifically is it reversible at low temperatures, i.e. can you prepare NF4ClO4 by mixing contacting cool NF3(g) and
FClO4(g)?
If so, then this would be very good because in apparently the same Christe paper also shows that fluorine perchlorate can be prepared by reaction
perchloric acid with fluorine, thus this too may be prepared as an anode gas by electrochemical fluorination, this time of KF.2HF melt containing
HClO4.
****
Putting all these points together, you might have a chance to form the desired NF4ClO4 in the second electrochemical fluorination step (where you
diffuse NF3(g) though your anode) if you have HClO4 in you KF.2HF melt and provided thermodynamics is not working against you in the reaction, else
your anode would merely preferentially form FClO4(g) over fluorinating NF3(g).
But I would hazard a guess that NF4+ wins thermodynamically because N-F bonds are possibly more favourable (stabler) than O-F bonds.
[Edited on 1-11-2013 by deltaH]
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woelen
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The problem with all this kind of chemistry is that no one (at least not of us) can test it. I would love to experiment with these chemicals and see
if I can make stuff like F-O-ClO3 and NF4(+) salts, but only the best equipped labs in the world can do this type of chemistry. And because it is of
academic interest only (I see no practical applications of the involved chemicals) I am afraid that it may take a long time before this kind of
research is done. I like to read about all the special things Christe has done, he really made remarkable compounds, but all of that is far beyond
what any home chemist can achieve.
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deltaH
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Sigh, you're right.
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