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Author: Subject: Ethyl Perchlorate
12AX7
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[*] posted on 20-7-2007 at 13:30


How reactive is trioxane? Seems to me it's an ether...

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[*] posted on 21-7-2007 at 01:55


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.

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[*] posted on 2-8-2007 at 10:48


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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[*] posted on 2-8-2007 at 13:57


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.

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[*] posted on 2-8-2007 at 20:25


"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. :)

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thumbup.gif posted on 3-8-2007 at 01:05


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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[*] posted on 3-8-2007 at 04:39


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.




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[*] posted on 3-8-2007 at 05:34


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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[*] posted on 3-8-2007 at 15:08


Oooh! Is there a [ClF6]ClO4? Pure oxidation! :D I don't want to think what that would do with powdered calcium metal... :D^2

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[*] posted on 4-8-2007 at 09:15


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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[*] posted on 11-8-2007 at 16:13


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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[*] posted on 12-8-2007 at 07:39


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.

.


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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[*] posted on 20-8-2007 at 07:05


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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[*] posted on 7-10-2007 at 20:52


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 1389 times

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[*] posted on 26-12-2008 at 12:53


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.




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[*] posted on 26-12-2008 at 15:15


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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[*] posted on 26-12-2008 at 19:47


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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[*] posted on 2-1-2009 at 07:31


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?



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[*] posted on 2-1-2009 at 12:34


Quote:
Originally posted by -=HeX=-
I read the paper, and the danger of death actually excites me.


Go to totse if that excites you :)

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[*] posted on 3-1-2009 at 05:59


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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[*] posted on 3-1-2009 at 06:24


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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[*] posted on 3-1-2009 at 06:40


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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[*] posted on 3-1-2009 at 09:26


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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[*] posted on 3-1-2009 at 17:32


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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[*] posted on 4-1-2009 at 10:51


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 :)




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