Yamato71
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What the hell did I just make?
Yes, it's THAT Yamato71 and I'm beginning to think I've got a learning disability.
TO WHIT:
While strolling down the aisle of my local Home Depot the other day to pick up some muriatic acid to clean my mineral specimens, I spied a bin full of
little white bags of pool shock for a buck each. Thinking that it was a great price for a pound of calcium hypochlorite, I tossed two into the
basket. It wasn't until I got home and read the label that I found out that I was the proud new owner of two pounds of sodium
dichloro-s-triazinetrione.
Mildly annoyed, I stared at the name and, reaching back into my long ago organic chemistry training, tried to imagine what such a molecule might look
like. Failing that, I just looked it up online.
Turns out, it's the sodium salt of 2,4-dichloroisocyanuric acid, a six-membered ring of alternating carbon and nitrogen atoms. The carbons at the 1,
3 and 5 positions are all bonded to oxygen (C1 to OH, C3 and C5 to keto groups). The nitrogens at positions 2 and 4 bond to chlorines, while the
nitrogen at position 6 has no substituent, but is instead double bonded to C1, making the molecule a lactam, and rendering the OH group next door on
C1 acidic though keto/enol tautomerization. The ring cannot be aromatic because of the keto groups on C3 and C5.
This little molecule is getting interesting. Apparently, it is a strong oxidizer, yet it appears that it can itself be oxidized under fairly mild
conditions. So, I decide to experiment. First, let's compare its oxidizing activity with its predecessor, calcium (or sodium) hypochlorite. We all
know what happens when Ca(OCl)2 is mixed with a polyol like glycerine or brake fluid. I put an ounce or so of the pool shock in a plastic cup, set it
on the ground and added enough plain jane brake fluid to saturate it, and waited...
and waited..., and waited a little more. Nada, zip, zilch. No fire, no smoke, heck, even the sides of the cup were cool to the touch. After about
ten minutes, I played a torch flame across the sticky mass and got an immediate reaction.
At first, the brake fluid just burned with an orange smoky flame. After a few seconds however the flame changed to bright yellow and the smoke went
away. Something was oxidizing the fluid and judging by the smell of free chlorine in the air, it was probably oxygen from the keto groups attacking
the brake fluid. The reaction never reached the level of violence I would have expected had I used a hypochlorite salt.
Ok, this stuff attacks organic material, but not as vigorously as hypochlorite. Let's see if we can entice the chlorine atoms to participate. Lets
see what happens when we try to reduce the pool shock with powdered metals, starting with zinc dust. This time, I ran the chunky oxidizer through a
coffee grinder to make for a more homogenous reaction mixture. I weighed out roughly equimolar amounts based on the chlorine content of the pool
shock, placed both powdered oxidizer and zinc dust in a paper cup, gently shook the cup until the color was a uniform grey, sat it on the ground and
stepped back a few meters.
After waiting a few minutes to see if the mixture would self ignite, which it didn't, I again touched a torch flame to the surface of the mixture. I
noticed that a small spot in the center was glowing dull red and that the glow was spreading slowly, but evenly, through the reaction mixture. The
most noticeable aspect of the reaction was it's gentle nature. It appeared to proceed with an essentially zero order rate. It never accelerated nor
slowed down, maintaining a reddish yellow glow at the reaction front. I remember thinking that this mixture would make a kick ass delay composition
when I noticed something else. The air for maybe fifty meters around me resembled a fog bank, yet very little smoke was coming from the reaction. Of
course! Anhydrous zinc chloride was being emitted as a gas from the hot reaction. The invisible gas would drift for a short distance before either
condensing back to an ultra fine powder, reacting with the humidity in the air, or both. That little pile of grey powder made one heck of a smoke
composition, just like the military zinc-hexachloroethane smoke mixtures that have been in use since WWI. I have to write this one down for future
research.
I tried one more reduction using paint grade aluminum. As expected, the reaction was somewhat more vigorous, but was still fairly tame. I elected
not to try magnesium. It was time to move on to oxidation.
For my first attempt to oxidize sodium dichloro-s-triazinetrione, I prepared an equal (w/w) mixture of finely powdered pool shock and potassium
nitrate. I didn't expect the reaction to be vigorous, and I wasn't to be disappointed. It took a few seconds of torch flame to get the mixture to
light off and when it did, it was very sluggish. A great deal of caustic smelling white smoke evolved and the stench of free chlorine was almost
overpowering. There was another familiar, yet faint aroma I picked up on, but couldn't quite place it. I probably should have tried a little harder
to remember where I had smelled it before because it might have prevented a close call.
Just as I had reasoned with the brake fluid experiment, if I could get that free chlorine to participate in the reaction, I might come up with another
great smoke composition. To that end, I replaced the potassium nitrate oxidizer with finely powdered, non-prilled ammonium nitrate, figuring that the
free chlorine would be sequestered by any ammonia released by the decomposition of the ammonium nitrate. Ammonium chloride would create smoke just as
had the zinc chloride of the previous experiment. Again, equimolar portions of pool shock and ammonium nitrate were prepared and brought to the spot
where the other experiments had been conducted. As before, I dumped one of the powdered reactants atop the other and prepared to gently shake them
until the mixture was homogenous. It was at this point that the experiment didn't behave as predicted.
As the two powders came together and I began to gently shake the cup I immediately sensed that something wasn't kosher. Instead of the powder grains
sliding across one another and mixing, they instantly clumped. Within a few seconds, the mixture was immobile, just a pasty clump in the cup that
appeared to be fusing and becoming more fluid by the second. I reasoned that maybe the NH4NO3 was absorbing humidity, but it didn't feel like that
all. Besides, it was a fine Autumn evening with very low humidity. Just as I was mulling over the situation, another sensation jolted me to
attention. The side of the paper cup was becoming warm to the touch. At that moment I had only one thought. Put. It. Down...... NOW!
That done, I turned heel and retreated to what seemed like a safe distance and waited. I didn't have to wait long. Maybe a minute after I set the
paper cup down I was startled by a powerful detonation. The clump that was about the size of a grape blasted the cup into confetti, leaving a divot 6
inches across and 4 inches deep in the ground. Most of the reaction mass was spattered across a radius of ten meters. Evidently, only a small
portion of the material actually detonated. I had just seen two powdered solids react to form a liquid product which somehow became its own solvent.
Again, I smelled a little chlorine, but this time the other smell was much stronger and I recognized it instantly, chloramine!
What do you think happened here? If I had to hazard a guess, (no pun intended), I'd say the mysterious liquid product was a mixture of the three
chloramines, including the bad ass of the family, nitrogen trichloride. Once it began to solvate the reactants, the reaction went into thermal
runaway, detonating the NCl3. The one absolute takeaway here is that I won't be repeating the experiment. I strongly recommend that none of you try
it either.
Peace/out
Y71
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Tdep
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Hello THAT Yamato71, what a terrifying experiment.
Yes i'd say nitrogen trichloride as the main suspect too. Ammonium salts and hypochlorite are known conditions to produce the trichloroamine IIRC, and
if you smelt the very unpleasant chloroamines, then that's a bit of a giveaway.
The other highly unstable species you possibly (maybe likely?) created is ammonium chlorate. That also has a nasty reputation for detonating
unexpectedly, and I can see it forming fairly readily in the mix given a bit of water in the ammonium nitrate and a bit of runaway to decompose the
hypochlorite (which would contain a small amount of chlorate in it to begin with)
Edit: Oh I slightly misread, the accident happened from ammonium nitrate and NaDCC not with hypochlorite. Yeah, ammonia and NaDCC will make NCl3. Free
ammonia and TCCA (and possibly NaDCC?) detonate on contact, given concentrated enough ammonia. Ignore that bit about ammonium chlorate then.
[Edited on 28-11-2016 by Tdep]
[Edited on 28-11-2016 by Tdep]
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j_sum1
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Wow. Several comments.
1. Nice to have you back in the game. I don't need to tell you to take care with anything energetic and potentially surprising.
2. Reminds me of my first accidental synthesis of chloramine. deltaH came up with a suggestion here. And my experimentation and discussion begins here. It can't have been more than a few miligrams in a small beaker but it went off a bit unexpectedly and gave me quite a startle.
3. Yes, Na-DCCA really is an interesting compound. There are a few threads around but most of them are a mixed bag of ideas and half-complete
experimentation. They are worth a read. woelen started on on this and TCCA some years back. You do have to trawl through a few threads to get all
the info. I would do some reading on woelen's site for more. But AFAIK, no one here has done a whole lot on the energetic properties of Na-DCCA.
4. Like Ca(OCl)2 and trichloroisocyanuric acid, NaDCCA can be used effectively as a chlorine generator. NaDCCA is reasonably soluble which could be
an advantage in some contexts.
5. Most of my experimentation in this area involves complexes. A beautiful purple complex is formed with Cu2+ ions. It is easy to make. I tried
making the potassium analogue with only partial success. The literature states that complexes can be formed swapping Cu2+ for Pb, La and a few other
transition metals. I had a student experiment a bit but I would not vouch for the results. (And I cannot remember where I found a useful paper that
listed a few compounds.)
6. Others have said this, but you really should write a book.
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Yamato71
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In looking at the ring structure of DCCA salts it seems likely that another hideously toxic compound could arise during its thermal composition. Look
carefully at any of the three ring carbons, C1, C3 or C5. All three are keto carbons linked to nitrogen atoms on either side. (C1 tautomerizes
between =O and -OH). Of course, that four atom fragment is urea, or a chloro-substituted urea in this case. If during decomposition, either or both
of the adjacent C-N bonds cleave, then re-form on a loose chlorine atom, the resulting product would be our old friend phosgene, COCl2, a useful
synthetic reagent and military nerve agent. Other probable molecular fragments would be nitrogen oxides (Nx-Oy), chlorine dioxide (ClO2), nitrosyl
chloride (NOCl), hydrogen cyanide (HCN), cyanogen chloride (Cl-CN), sodium cyanide (NaCN), cyanogen (NC-CN), carbon monoxide (CO) and one of the
greatest killers of all time, sodium chloride (NaCl). 
Yet, as toxic as the products of runaway thermal decomposition of DCCA and its salts might be, I feel that under controlled conditions cyanuric acid
and its halogenated derivatives might be useful in the synthesis of any number of useful intermediates, including the compounds listed above.
Provided that no organic contamination is present, pyrolysis and fractional distillation of DCCA under reduced pressure or an inert atmosphere could
make some hard to get precursors OTC with a little effort.
[Edited on 28-11-2016 by Yamato71]
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woelen
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Na-DCCA and TCCA both react with ammonium ions and with ammonia explosively violently. These reactions really are surprising.
Just for fun, put a small amount of Na-DCCA or TCCA (not more than 1 gram and do this outside) on a dish and then pour a few ml of 25% ammonia on it.
You will be shocked by the violence of this reaction. Even 12% ammonia gives a scaringly violent reaction.
Best is to pour the ammonia from a construction with a stick of 1 meter length, so that you are not too close to the reaction mixture!
I am glad to read that you did not get an accident with your mixture. It could have been very different. I hope you learnt something from this: never
work with such large quantities if you don't know what can happen. Work with 100's of milligrams, one gram at most and only scale up if you know what
happens in the microscale experiment.
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Yamato71
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@ j_sum1 I just finished reading your posts detailing your efforts to synthesize anhydrous NO2 with urea and calcium hypochlorite. For someone who
was experimenting "in the blind" as I was tonight, your work proceeded in a very cautious and logical manner and you managed to resist the urge to
scale up or rush your experiments. All of us have, at one time or another, given in to that "more, bigger, faster" urge to save time. Most of us
have gotten away with it. Some of us didn't. Anyway, I just wanted to point out what you were doing right. Much respect sir.
Now, a few tidbits I'd like to add regarding the synthesis you attempted.
In the first phase you were reacting calcium hypochlorite and urea in the presence of slaked lime. You mentioned that a popping began as you were
mixing the dry reagents. That sounds almost exactly like what happened to me tonight. I believe, in your case, you forgot that your oxidizing agent
was a powerful chlorinating agent as well. Exposing Ca(OCl)2 to either harsh acidic or basic conditions will result in the creation of free chlorine
and/or monoatomic chlorine radcals which would have made short work of the two amine groups on the urea molecule. As soon as each nitrogen added a
third chlorine, its covalent bond with the carbonyl carbon would vanish. Almost immediately two things would occur. The newly formed molecule of
nitrogen trichloride would drift away from the carbon and a chlorine radical would take its place. The same radical halogenation would occur at the
other nitrogen, completely replacing both amine groups with chlorine atoms. Instead of solid urea, you now have gaseous phosgene, an insidious poison
which leaves the reaction mixture, driving the reaction to the right, increasing the rate of phosgene production. Meanwhile, the NCl3 that is
building up in the reaction mixture suddenly notices that it's really hot and the pH is unbearably high. It isn't happy. Any time the local
concentration of NCl3 exceeds the critical threshold, it makes its displeasure known to anyone within earshot.
I'm afraid that urea is just about the worst possible reactant you can use if your goal is make NO2. NO2 reacts vigorously with urea. All the
products of this reaction are colorless gases, nitrogen, water vapor, CO and CO2 which, as in the previous example, leave the reaction.
LeChattelier's principal kicks in and the reaction rate kicks into high gear in an attempt to make more product and reestablish equilibrium between
the forward and reverse reactions. If that weren't bad enough, the water produced by the reaction acts as a catalyst, speeding up the destruction of
NO2. Dropping urea prills into red fuming nitric acid is an easy way to convert it into white fuming nitric acid as urea destroys dissolved NO2 but
does not react with HNO3. When the fizzing stops, you're done.
This reaction made world news last year when it was exploited by Volkswagen in a recent US emissions testing scandal. The level of oxides of nitrogen
(NOX) at the tail pipe had to meet US EPA standards before the company's diesel cars could be imported. Knowing that they couldn't pass the tests, VW
did what anybody faced with the same situation would do. They cheated. Canisters of dry prilled urea were hidden in the exhaust system of the test
vehicles. NOX entering the canisters exited as water vapor, nitrogen and CO2. Since the tests were specific for NOX, the kapooter ignored the
abnormal CO2 emissions. It wasn't until a human being correlated the low NOX/high CO2 pattern that anybody suspected that a respected multi billion
dollar world corporation was trying to slip its pecker into the EPA's pocket. The issue that really pushed the bizarre meter up past eleven was the
fact that the guilty scumbag corporation wasn't even Chinese, it was German.
Anyway, I can think of maybe a dozen ways to make dry NO2 from OTC ingredients other than urea. I assume that air is OTC enough for you? When I get
a chance, I'll post some of them to the other thread.
Snap, Crackle, BANG!
Y71
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Morgan
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Someone posted this before, but perhaps related. In the last example he adds NH4NO3 to the NH4OH.
https://www.youtube.com/watch?v=9A9Fg-hJy-4
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Cryolite.
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Na-DCCA is the mono sodium salt of dichlorocyanuric acid. This dichlorocyanuric acid is simply the double chloroimide of cyanuric acid, which is
industrially prepared by bubbling chlorine gas into alkaline cyanuric acid solutions. The N-Cl bonds are very labile, and so Na-DCCA can almost be
thought of as a nonacidic source of Cl+ in solution. Once these chlorines have been reacted off and H+ is allowed to take their place, however, the
remaining cyanuric acid is effectively inert and usually drops out of the reaction mixture.
Cyanuric acid, incidentally, is just the cyclic trimer of hydrogen cyanate, which easily decomposes to carbon dioxide and ammonia in contact with
water. The decomposition of trisodium cyanurate to sodium cyanate, followed by the reduction with activated charcoal, forms the basis to produce
sodium cyanide.
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j_sum1
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Quote: Originally posted by Yamato71  | @ j_sum1 I just finished reading your posts detailing your efforts to synthesize anhydrous NO2 with urea and calcium hypochlorite. For someone who
was experimenting "in the blind" as I was tonight, your work proceeded in a very cautious and logical manner and you managed to resist the urge to
scale up or rush your experiments. All of us have, at one time or another, given in to that "more, bigger, faster" urge to save time. Most of us
have gotten away with it. Some of us didn't. Anyway, I just wanted to point out what you were doing right. Much respect sir.
Now, a few tidbits I'd like to add regarding the synthesis you attempted.
In the first phase you were reacting calcium hypochlorite and urea in the presence of slaked lime. You mentioned that a popping began as you were
mixing the dry reagents. That sounds almost exactly like what happened to me tonight. I believe, in your case, you forgot that your oxidizing agent
was a powerful chlorinating agent as well. Exposing Ca(OCl)2 to either harsh acidic or basic conditions will result in the creation of free chlorine
and/or monoatomic chlorine radcals which would have made short work of the two amine groups on the urea molecule. As soon as each nitrogen added a
third chlorine, its covalent bond with the carbonyl carbon would vanish. Almost immediately two things would occur. The newly formed molecule of
nitrogen trichloride would drift away from the carbon and a chlorine radical would take its place. The same radical halogenation would occur at the
other nitrogen, completely replacing both amine groups with chlorine atoms. Instead of solid urea, you now have gaseous phosgene, an insidious poison
which leaves the reaction mixture, driving the reaction to the right, increasing the rate of phosgene production. Meanwhile, the NCl3 that is
building up in the reaction mixture suddenly notices that it's really hot and the pH is unbearably high. It isn't happy. Any time the local
concentration of NCl3 exceeds the critical threshold, it makes its displeasure known to anyone within earshot.
I'm afraid that urea is just about the worst possible reactant you can use if your goal is make NO2. NO2 reacts vigorously with urea. All the
products of this reaction are colorless gases, nitrogen, water vapor, CO and CO2 which, as in the previous example, leave the reaction.
LeChattelier's principal kicks in and the reaction rate kicks into high gear in an attempt to make more product and reestablish equilibrium between
the forward and reverse reactions. If that weren't bad enough, the water produced by the reaction acts as a catalyst, speeding up the destruction of
NO2. Dropping urea prills into red fuming nitric acid is an easy way to convert it into white fuming nitric acid as urea destroys dissolved NO2 but
does not react with HNO3. When the fizzing stops, you're done.
This reaction made world news last year when it was exploited by Volkswagen in a recent US emissions testing scandal. The level of oxides of nitrogen
(NOX) at the tail pipe had to meet US EPA standards before the company's diesel cars could be imported. Knowing that they couldn't pass the tests, VW
did what anybody faced with the same situation would do. They cheated. Canisters of dry prilled urea were hidden in the exhaust system of the test
vehicles. NOX entering the canisters exited as water vapor, nitrogen and CO2. Since the tests were specific for NOX, the kapooter ignored the
abnormal CO2 emissions. It wasn't until a human being correlated the low NOX/high CO2 pattern that anybody suspected that a respected multi billion
dollar world corporation was trying to slip its pecker into the EPA's pocket. The issue that really pushed the bizarre meter up past eleven was the
fact that the guilty scumbag corporation wasn't even Chinese, it was German.
Anyway, I can think of maybe a dozen ways to make dry NO2 from OTC ingredients other than urea. I assume that air is OTC enough for you? When I get
a chance, I'll post some of them to the other thread.
Snap, Crackle, BANG!
Y71 |
Thanks Yamato71. That's kind words.
Short answer, project over. I now realise that the direction of my experimentation was a dead end as far as my aim is concerned. But at the time I
was a newbie and really did not appreciate the difference in the level of oxidation between urea and nitrates. I do now and it is really good to see
how far I have come in a couple of years. And, as you say, between the carbonyl groups, the highly oxidising conditions and the wandering chlorines
(in a +1 oxidation state) there is a lot of potential for side reactions and not really any chance of producing NO2.
Man that thread was a good learning curve for me. Thanks again aga.
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Bert
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Considering the condition of your hand, lucky you could feel the self heating in time to get rid of that dixie cup-O-mayhem.
How is the grass around the site behaving?
Rapopart’s Rules for critical commentary:
1. Attempt to re-express your target’s position so clearly, vividly and fairly that your target says: “Thanks, I wish I’d thought of putting it
that way.”
2. List any points of agreement (especially if they are not matters of general or widespread agreement).
3. Mention anything you have learned from your target.
4. Only then are you permitted to say so much as a word of rebuttal or criticism.
Anatol Rapoport was a Russian-born American mathematical psychologist (1911-2007).
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cagse3888
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Yamoto, I had a similar experience with trichloroisocyanuric acid, KNO3, and kerosene. I was trying to make a STP cast rocket motor. I ended up with a
much bigger bang than expected.
Now, I ended up actually using the pool counterpart, spa brominating tablets, bromochlorodimethylhydantoin is much more stable in these types of
mixtures.
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Bert
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| Quote: Originally posted by cagse3888 | | Yamoto, I had a similar experience with trichloroisocyanuric acid, KNO3, and kerosene. I was trying to make a STP cast rocket motor.
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Please elucidate.
Rapopart’s Rules for critical commentary:
1. Attempt to re-express your target’s position so clearly, vividly and fairly that your target says: “Thanks, I wish I’d thought of putting it
that way.”
2. List any points of agreement (especially if they are not matters of general or widespread agreement).
3. Mention anything you have learned from your target.
4. Only then are you permitted to say so much as a word of rebuttal or criticism.
Anatol Rapoport was a Russian-born American mathematical psychologist (1911-2007).
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Metacelsus
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That sounds like a recipe for an unstable pipe bomb, not a rocket motor.
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DubaiAmateurRocketry
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Hey I am just glad that youre back on the forum.
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Bert
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https://patentimages.storage.googleapis.com/pages/US3465674-...
Yep, kinda does.
Rapopart’s Rules for critical commentary:
1. Attempt to re-express your target’s position so clearly, vividly and fairly that your target says: “Thanks, I wish I’d thought of putting it
that way.”
2. List any points of agreement (especially if they are not matters of general or widespread agreement).
3. Mention anything you have learned from your target.
4. Only then are you permitted to say so much as a word of rebuttal or criticism.
Anatol Rapoport was a Russian-born American mathematical psychologist (1911-2007).
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cagce88
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Well Bert and Metacelsus, the previous mixture makes sense now. I was trying to make a castable rocket motor from chemical ingredients you could buy
in a farm store or hardware store. In addition, the composition could be mixed at room temperature and poured as a slurry.
Like i said, I tried a potassium nitrate/kerosene mix, but the two aren't exactly soluble together. I added the chlorine pool shock, and ended up with
a bang. This experiment was three years ago, and I remember it quickly blowing after 30 seconds. Then I tried spa-brominating tablets, which works
wonderfully. Also, a little bit of petro jelly mixed in with the kerosene will thicken the mixture to ensure the kerosene doesn't exude.
Is there better, safer ways to do what I was trying to do earlier, yes, probably. I'm all for easier and safer.
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Bert
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You want to make quick, simple, rocket engines from over the counter materials? And not have too many surprises???
Well, sounds like you already know where to find Potassium nitrate, so- Look here:
http://www.nakka-rocketry.net/pvcmot4.html
If you can't find sorbitol in your area?
http://m.ebay.com/sch/i.html?isRefine=true&_pgn=1&_n...
Rather better ISP than black powder, and they hardly ever set themselves off for no apparent reason.
[Edited on 8-12-2016 by Bert]
Rapopart’s Rules for critical commentary:
1. Attempt to re-express your target’s position so clearly, vividly and fairly that your target says: “Thanks, I wish I’d thought of putting it
that way.”
2. List any points of agreement (especially if they are not matters of general or widespread agreement).
3. Mention anything you have learned from your target.
4. Only then are you permitted to say so much as a word of rebuttal or criticism.
Anatol Rapoport was a Russian-born American mathematical psychologist (1911-2007).
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DraconicAcid
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Whenever the question is "What the hell did I just make?", the answer is probably, "A mistake."
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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j_sum1
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Yeah, but not always. This stuff has properties way beyond a truck deck liner.
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Herr Haber
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Ethanol ? (This was a while back, I agree)
Nitrocellulose ?
Teflon ?
CFC gasses ? (ok, those had unforseable side effects...)
I'm certain you can think of many more "mistakes"
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mayko
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I think the discovery of mauveine from a failed run at quinine is my favorite story in that genre. It even includes covert home labs!
| Quote: |
During the Easter vacation in 1856, while Hofmann was visiting his native East End, Perkin performed some further experiments in the crude laboratory
in his apartment on the top floor of his home in Cable Street in east London. It was here that he made his great accidental discovery: that aniline
could be partly transformed into a crude mixture which when extracted with alcohol produced a substance with an intense purple colour.[3] Perkin, who
had an interest in painting and photography, immediately became enthusiastic about this result and carried out further trials with his friend Arthur
Church and his brother Thomas. Since these experiments were not part of the work on quinine which had been assigned to Perkin, the trio carried them
out in a hut in Perkin's garden, so as to keep them secret from Hofmann.
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https://en.wikipedia.org/wiki/William_Henry_Perkin#The_accid...
al-khemie is not a terrorist organization
"Chemicals, chemicals... I need chemicals!" - George Hayduke
"Wubbalubba dub-dub!" - Rick Sanchez
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mysteriusbhoice
Hazard to Others
Posts: 477
Registered: 27-1-2016
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Mood: Became chemistry catboy Vtuber Nyaa
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if you mix ammonium nitrate with even calcium hypochlorite it will detonate but only in small pops!!
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BNCP
Harmless
Posts: 1
Registered: 5-8-2005
Member Is Offline
Mood: No Mood
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This is indeed very complex chemistry with partially explosive compounds. Without professional analysis it's just possible to give theoretical
explanations. Due to my experience in lab during the last 20 years the theory might be way off the reality sometimes... the potentially unstable
product should be NCl3 with the whole series of chlorite or chlorate salts. This might be autocatalytic to some extent where i would suspect the NCL3
to be the critical compound that caused the selfignition/explosion or detonation as a threshold level was reached. The analysis of such unstable
mixtures is a real challenge. There is a nice scientific article about a (irony on) "professional" (irony off) accident i stumbled upon:
http://onlinelibrary.wiley.com/doi/10.1002/prs.11741/pdf
It scares the shit out of me watching youtube videos of people synthesizing NCL3 in glass beakers or glass equipment without protective shield and
holding this shit in their hands. This is careless and stupid and will for sure sooner or later cause injury or loss of limbs or eyesight. My last
project in the company lab was about synthesizing a highly fluorinated molecule with a tetrazole side chain. Nasty reagents like TMS-N3, potential for
liberation of HF or hydrazoic acid, potentially ames positive compounds... Yamato i must admit that you might be right with the statement that you
have a learning disability, why do you hold such unstable and unpredictable mixtures in your hands? Even a few inches or cm of distance make a big
difference in blast injuries if you don't use 100g amounts... Darwinism is not just a theory ;-)
[Edited on 2-1-2017 by BNCP]
[Edited on 2-1-2017 by BNCP]
[Edited on 2-1-2017 by BNCP]
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