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12AX7
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How reactive is trioxane? Seems to me it's an ether...
Tim
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sparkgap
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Trioxane ain't just any ether; it's a trimer of formaldehyde.  franklyn seems to
be using it as a formaldehyde source in the scheme he outlined in his post.
sparky (~_~)
"What's UTFSE? I keep hearing about it, but I can't be arsed to search for the answer..."
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Sciocrat
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Can someone explain to me what makes ethyl perchlorate so much more powerful than other explosives?
And in general, what is the reason that some explosives, with very high potential energies in their chemical bonds, remain stable at room temperature,
and need a detonator to start the explosive reaction, even though every system in the nature tries to have as low potential energy as possible?
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12AX7
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Lots of systems require activation energy, even if the result if favorable. A boulder sitting precariously on the edge of a cliff would rather be at
the bottom, but it needs some initial push to do it. Same thing with molecules, the boulders are just smaller and closer to their cliffs.
Some systems do not require activation energy. (Obviously, spontaneous reactions don't.) Quantum mechanics allows tunneling, with probability over
time combining with activation energy. To fuse two hydrogen nuclei, some energy must be put in (despite the energy resulting from fusion into
helium), but between tunneling and thermal distributions, you can get by with a far lower average temperature, you just get a much slower rate. (The
probability at room temperature is finite -- extremely improbable, perhaps never having happened in the history of the universe, but probable
nonetheless!) Molecular systems are of course on a quantum scale, but being somewhat large assemblies, can be seen to work from a more mechanical
(classical) view. Molecular bonds are springy, so you could perhaps imagine atoms as balls covered in stretchy velcro. Some like to stick together
more than others, but many will stick together anyway, allowing for the precarious placement of nitrate groups near hydrogen and carbon.
Perchlorates (the salts) are, in general, more or less stable. The group is happiest as an anion, but an organic substance doesn't really make a good
cation, so the perchlorate group is grumpy and unstable instead. In terms of bonding, the oxygens hanging out with the chlorine would much rather be
hanging out with the carbon and hydrogen on the organic. By giving the group a negative (anionic) charge, it's much happier. Oxygen really loves
electrons -- that's why there are so many oxide and oxoanion (sulfate, silicate, etc.) minerals. There's four oxygens on the perchlorate ion, so even
with seven stolen from the chlorine center, they want one more.
More stable organic bases, amines for instance, make better salts. Ammonium perchlorate is pretty stable, though it can be detonated, as can ammonium
nitrate. Does anyone know about, say, methylamine perchlorate? 2 CH3NH3ClO4 --> 2 CO + 6 H2O + N2 + Cl2 looks very balanced.
Tim
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sparkgap
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"Some systems do not require activation energy. (Obviously, spontaneous reactions don't.)"
That is fallacious. All reactions need a "push", as you put it. Some are easier to push than others. 
sparky (^_^)
"What's UTFSE? I keep hearing about it, but I can't be arsed to search for the answer..."
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Sciocrat
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Tnx for the explanation, 12AX7.
| Quote: | Originally posted by sparkgap
"Some systems do not require activation energy. (Obviously, spontaneous reactions don't.)"
That is fallacious. All reactions need a "push", as you put it. Some are easier to push than others.
sparky (^_^) |
You mean like 2H2 + O2 -> 2H2O?
It is a spontaneous reaction, and it happens at room temperature, but very slowly. So it needs that "push", to allow the reaction to happen at a much
greater speed.
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vulture
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A review of the synthesis and reactions of organic perchlorates is given in Russ. Chem. Rev. 1988, 57, 1041.
Unfortunately I am unable to acces this article.
One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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JohnWW
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Further to 12AX7, I would note that the stability of ionic perchlorates, containing the ClO4- anion, is due to its being tetrahedrally symmetric, and
resonance-stabilized so that the negative charge and associated single Cl-O bond are distributed equally over all 4 O atoms. The symmetry of the
anion, and also the octahedral symmetry of the recently-discovered cation ClF6+ (made by the reaction of [KrF]SbF6 with ClF5, which releases Kr gas,
and a F+ cation which combines electrophilically with the ClF5), makes the +7 oxidation state the most stable of the positive oxidation states of
chlorine. This symmetry also accounts for ClO4- being much more stable than ClO3-, which has an unutilized electron pair.
By comparison, covalent esters of perchloric acid (one of the strongest known acids) do not have this energy-lowering symmetry and
resonance-stabilization, resulting in detonation of such esters being more exothermic than detonation of a mixture of an alcohol and an ionic
perchlorate.
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12AX7
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Oooh! Is there a [ClF6]ClO4? Pure oxidation!  I don't want to think what
that would do with powdered calcium metal... ^2
Tim
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JohnWW
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That theoretically could be made, by adding perchlorate to the reaction solution (which I think is in either liquid HF or a fluorocarbon) afterwards,
and evaporating it down to crystallize either the [ClF6]ClO4 or the hexafluoroantimonate, whichever is least soluble.
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franklyn
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| Quote: | Originally posted by 12AX7
Does anyone know about, say, methylamine perchlorate?
2 CH3NH3ClO4 --> 2 CO + 6 H2O + N2 + Cl2 looks very balanced. |
Methylamine Perchlorate has been known a long tiime, here's a patent description posted by
Rosco Bodine
http://www.sciencemadness.org/talk/viewthread.php?action=att...
Triaminobenzene =>
http://cdb.ics.uci.edu/CHEMDB/Web/cgibin/ChemicalDetailWeb.p...
is tribasic and is commercially available as Triaminobenzene Trihydrochloride. Added to Sodium
Perchlorate in water will precipitate NaCl leaving mostly a TriperchloroTriaminobenzene solution.
Not sure how to partition those with solvent though.
.
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not_important
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| Quote: | Originally posted by franklyn
...
Triaminobenzene =>
http://cdb.ics.uci.edu/CHEMDB/Web/cgibin/ChemicalDetailWeb.p...
is tribasic and is commercially available as Triaminobenzene Trihydrochloride. Added to Sodium
Perchlorate in water will precipitate NaCl leaving mostly a TriperchloroTriaminobenzene solution.
Not sure how to partition those with solvent though.
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AgClO4 or NaClO4 in MeOH, EtOH, or possibly acetone, as solvent; NaCl has rather low solubility in those solvents. will AgCl is almost zero.
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tito-o-mac
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So does nitrosyl perchlorate requirethe same principal and steps to synthesise like ethyl perchlorate , or like the other family of perchlorates.
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Axt
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No, nitrosyl perchlorate is most readily formed by N2O3 on HClO4. You got no reply because it has its own thread : http://www.sciencemadness.org/talk/viewthread.php?tid=196
Here an article on the perchlorate esters of ethylene glycol, glycerine and pentaerythritol.
Attachment: perchlorate esters of polyhydric alcohols.pdf (479kB)
This file has been downloaded 1495 times
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-=HeX=-
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I apologise for the necromancy... But I have been doing a LOT of theoretical work on Glycerin TriPerchlorate. Glycerin TriPerchlorate is the
perchloric ester of Glycerin, and a direct relative of NitroGlycerin. Think NG with ClO3+ replacing the NO2+ groups. I cant quite explain, so I will
eventually draw it out. Anyways, I have one problem and thats its decomp reaction. There are 2 possibilities. Here they are:
2C3H5(OClO3)3 ==> 6CO2 + 3Cl2 + 5 H2O + 3.5O2
OR IS IT...
2C3H5(OClO3)3 ==> 6CO2 + 6HCl + 2H2O + 5O2
The problem is... I can never remember whether oxygen or chlorine has the greater affinity for hydrogen... Also, is the ClO3+ ion stable enough to
survive even the esterification?
I am assuming it will be touch sensitive at least, hydroscopic, and insanely powerful. When I get around to dealing with it, I will tell you if I
survive.
If you give a man a match he will be warm for a moment. Set him alight and he will be warm for the rest of his life.
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kclo4
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I know that HCl can be oxidized to produce H2O and Cl2 using Oxygen. Oxygen is more electronegative I am pretty sure.
Although, I'm sure the structure of the molecule is going to make a difference on its decomposition products.
Don't you mean the ClO4- Ion?
ClO3+ doesn't exist, but ClO3- is very unstable isn't it?
I don't think it would be wise to make, although I haven't studied it at all, it just seems to dangerous.
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Formatik
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| Quote: | Originally posted by -=HeX=-
I apologise for the necromancy... But I have been doing a LOT of theoretical work on Glycerin TriPerchlorate. Glycerin TriPerchlorate is the
perchloric ester of Glycerin, and a direct relative of NitroGlycerin. Think NG with ClO3+ replacing the NO2+ groups. I cant quite explain, so I will
eventually draw it out. Anyways, I have one problem and thats its decomp reaction. There are 2 possibilities. Here they are:
2C3H5(OClO3)3 ==> 6CO2 + 3Cl2 + 5 H2O + 3.5O2
OR IS IT...
2C3H5(OClO3)3 ==> 6CO2 + 6HCl + 2H2O + 5O2
The problem is... I can never remember whether oxygen or chlorine has the greater affinity for hydrogen... |
I think the same as above. O2 is more electronegative than Cl2, so it could be the first reaction.
| Quote: | | I am assuming it will be touch sensitive at least, hydroscopic, and insanely powerful. When I get around to dealing with it, I will tell you if I
survive. |
The paper by Axt right above your post mentions glycerol perchlorate ester. They also made pentaerythritol and glycol esters. The procedures and
work-up are quite dangerous. Without a solvent and handling the raw esters, they explode even on simple decantation. Glycol perchlorate is already
stronger than NG but it is also so sensitive, that it explodes when water is added to it. It is known glycols as well as glycol ethers mixed with with
even 70% HClO4 decomposes violently (proceeding through esters) at regular temperatures. The reason they also work at dry ice temperatures in the
paper above in the beginning phase of reaction, but they are also using anhydrous HClO4 for esterification, which itself is already violently more
reactive than the azeotropic acid.
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-=HeX=-
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I read the paper, and the danger of death actually excites me. I am going to work more on the theory while I rebuild my lab, and then consider
attempting the synth. Chemicals are never a problem for me, as I have several good sources of them. Would a H2SO4 (w/ SO3) and 70% HClO4 mix work
instead of 100% HClO4?
If you give a man a match he will be warm for a moment. Set him alight and he will be warm for the rest of his life.
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kclo4
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| Quote: | Originally posted by -=HeX=-
I read the paper, and the danger of death actually excites me. |
Go to totse if that excites you
*smacks in face with a tightly rolled newspaper*
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497
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| Quote: |
and the danger of death actually excites me. |
Hey, at least he's doing it with style. I can respect that. Not some piece of crap BP-in-metal-pipe that's going to kill him, it'll be real live
triperchlorate ester of glycerol, now is that classy or what?
I wonder how useful the partially perchlorated polyhydritic alcohols could be? Blaster synthesised one apparently, and said it was fairly stable,
maybe something like like mono/di/tri/tetra/pentaperchloratomannitol would be usable?
Hell, while you're at it, why not mix in some other fun energetic groups? Maybe triazidotriperchloratomannitol or dinitrodiperchloratoerytritol?
As far as perchlorate salts go, what about the perchlorate salt of TATB? Is that even possible? Or maybe the perchlorate salt of
triazidotriaminobenzene?
It seems there are so many possibilities and combinations that one of them must be a practical explosive... The trick would be finding it without
getting killed first by all the ones that aren't..
[Edited on 3-1-2009 by 497]
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garage chemist
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For esterifying polyhydric alcohols with perchloric acid, I imagine a mix with H2SO4/SO3 would be unsuitable- remember, HClO4 is a stronger acid than
H2SO4, so the mix would preferentially form sulfuric esters. Also, SO3 is too aggressive towards oxidisable organics. It is a harsh oxidant that likes
to turn organics into black gunk.
Instead, learn about perchloric anhydride, Cl2O7, and its preparation from anydrous HClO4 and P2O5. This is actually less exlosive than the anhydrous
HClO4 used to prepare it.
[Edited on 3-1-2009 by garage chemist]
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497
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How powerful of a dehydrator does it take to make Cl2O7? It would be nice if others such as B2O3 or HPO3 could be used instead of P2O5..
This got me wondering, is there a reference out there that has some kind of ranked list of dehydrators? Like P2O5>B2O3>SO3>N2O5>etc? The
only thing like it that I've seen was in reference to desiccants, so obviously most of the interesting compounds were left out. It would be nice to
know what will dehydrate what...
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woelen
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Such a list is useless, because drying agents also have other properties.
E.g. SO3 is a VERY good drying agent, but it also is strongly oxidizing and turns many organics into black tarry crap.
As another example: SOCl2 also is an extremely good drying agent, but it also is a decent reductor and is capable of replacing -OH groups by -Cl in
many organics and even some inorganic compounds.
For each specific application you have to carefully select a drying agent.
[Edited on 3-1-09 by woelen]
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497
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I didn't really have organic compounds in mind. I was more thinking about which inorganic compounds will dehydrate other inorganics. Like (HPO3)n +
H2SO4 -> H3PO4 + SO3.
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-=HeX=-
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Nice... Telling me to go to totse *sighs* I am not compatible with k3wl scum. What I meant was that if it does blow me sky high, I will have died with
class. Now if I survive it, even better I am already starting the preparations
of fixing blast screens, getting protective gear, installing a fume hood of sorts, etc. I am, after all, going all out on this, and wish to maybe
improvise it if the school reneges on the deal.
I may look at mixed group ones, like glycerol diperchlorate azide or something like that, and may find one with acceptable sensitivity. The maths
works, but will the reaction? thats the million euro question. I an making drop test rigs, etc and wish to prepare perchloric esters of glycerin,
mannitol, erythritol, pentaerythritol, and anything else I can think of. You never know, I may come back alive I could have something nice
If you give a man a match he will be warm for a moment. Set him alight and he will be warm for the rest of his life.
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