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Author: Subject: TCCA, Na-DCCA and cyanuric acid
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[*] posted on 3-5-2006 at 10:14


Organikum is right with the warnings. The reaction is really exothermic and not cooling it with ice will make you sorry for sure. If cooled and if only a minimum of H2SO4 is used as catalyst then it proceeds slowly and safely (1 drop of H2SO4 per 500ml of acetone is OK).
See also the one of the last posts in Question on making CH3NO2? where I gave a few more advices on the synthesis.

The Good Read: Trichloroisocyanuric Acid - A Safe and Efficient Oxidant should also be mentioned here.

For Woelen:

[Edited on 3-5-2006 by Nicodem]

Attachment: trichloroisocyanuric_review.pdf (148kB)
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[*] posted on 3-5-2006 at 11:28


Quote:
Is your TCCA scented?

No, it is not scented. It is almost 100% pure. No gas with H2SO4 (so, no carbonate fillers), no smell, except the "swimming pool" chlorous smell.

Good that you warned me with the TCCA/acetone mix. I was tempted to heat the solution of TCCA in acetone and that could give bad surprises :o. Strange that acetone and TCCA alone do not react, but that a trace of acid apparently gives a very exothermic reaction. I'll certainly try it (on a very small scale to start with). But for now, I first am going to do my homework: "TClCA.rar" :) (thanks Nicodem !!).

I, however, could not download the file Trichloroisocyanuric Acid - A Safe and Efficient Oxidant.pdf . :( Is the link broken?

EDIT: I tried the reaction. I took 1.5 ml of acetone and added a granule of TCCA (appr. 4 mm diameter, a small distorted cube). No reaction could be observed, and no heat was evolved. Next, I added a single drop of conc. HCl which I let flow into the liquid by applying it along the glass. The liquid turned light green and then white and turbid. The small piece of TCCA still dissolves as well as without the acid, but now considerable heat is produced, and the liquid becomes turbid. When the entire piece of TCCA is dissolved, the liquid is quite hot. So, the reaction was not spectacular to see, but it was very instructive to feel how much heat is produced by such a small piece of TCCA, with all the heat spread out over the full 1.5 ml of acetone. When I scale up the experiment, I will take your warnings into account seriously.

[Edited on 3-5-06 by woelen]




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[*] posted on 4-5-2006 at 06:01


Quote:

If cooled and if only a minimum of H2SO4 is used as catalyst then it proceeds slowly and safely
Not really. If not enough acid is used to start the reaction it just doesn´t start, but as soon it kicks in it never preceeds slowly and safely but always is a chainreaction prone to runaways.

The safe way to do it is to add the TCCA in portions to a cooled mixture (10°C) of water/acetone and some acid. If after the first portion of TCCA no reaction kicks obviously in then add more acid until it does. NEVER add more TCCA or heat then or you might regret it. I did.
This may take some good time, I just added the TCCA every 45 minutes or so over several hours. Better safe than sorry.

Look here

/ORG




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[*] posted on 4-5-2006 at 11:17


I did the chlorination of acetone only twice but I think it was enough to learn a lesson. The first time I thought to add diluted H2SO4 as a catalyst in order to measure more accurately the minute amount needed. The reaction did not start, so I added more and more. I even removed the ice bath, stupidly. After I added almost 1ml of 35% H2SO4 in total, to the mixture of 140ml acetone and 42g TCCA, the reaction started all of a sudden and the whole thing went into a violent reflux that my condenser barely managed. I dare not to think what would have happened if the chloroacetone fumes escaped :(!
The second time I followed the French patent more accurately and used 96% H2SO4 instead. Into 1L flask in an ice bath I added 440ml of acetone under intensive magnetic stirring. I added several drops of 96% H2SO4 (less than 0.5ml) and trough an addition funnel slowly added a solution of 121g TCCA (0.52 mol) in 360ml acetone. The flask had a reflux condenser for safety in case of a runaway reaction and a thermometer. The reaction started immediately after starting the addition of the TCCA solution as indicated by a temperature rising. After ending the addition the temperature rose very slowly up to 40°C and melted most of the ice bath in the course of about two hours. Meanwhile the cyanuric acid begun to precipitate. The sulfuric acid was then neutralized by adding about 1g Na2CO3, then the precipitate vacuum filtered and washed with 50ml acetone and the filtrate distilled. The fraction distilling above 90°C was collected and redistilled, collecting the fraction distilling at 110-120°C. There was collected approximately 100ml of the nasty product (I didn’t even measure or weight it accurately enough just to avoid too much contact).

I think that if one uses concentrated H2SO4 the reaction does not have an induction period and starts immediately while the reaction speed can be regulated by the amount of the acid. This is also consistent with theory since conc. H2SO4 easily enolize the acetone while the diluted one is nearly not as efficient (water is about 8 magnitudes more “basic” than acetone and thus almost all the protons are “neutralized” by H2O). The diluted acid still very strongly activates TCCA (makes it electrophilic) but there is not enough enolized acetone around to maintain the reaction. At least this is my interpretation on why this happens.




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[*] posted on 4-5-2006 at 13:35


Quote:

I think that if one uses concentrated H2SO4 the reaction does not have an induction period and starts immediately while the reaction speed can be regulated by the amount of the acid.
Well I thought you better regulate it by using enough acid to make start for sure and then by further addition of the TCCA.

I also preferred in the longer run to add the TCCA as solid and not dissolved in acetone for TCCA dissolved in acetone is an accident waiting to happen. Whatever kicks the chainreaction in, some impurity or some sunlight. And you have a lots of trouble.

But of course this deals with anhydrous acetone + TCCA IIRC, what happens in the end in an aqueous solution might be completely different.




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[*] posted on 5-5-2006 at 14:19
Homebrew primary alcohol oxidation


Well, I have Acetonitrile but I paid $123 for 4 L and it is a deadly poison so I cheaped out and used acetone instead. I carefully dissolved TCICA in acetone. DId not weigh the TCiCA but I had roughly 300 ml of acetone. I would estimate I only dissolved maybe 20 grams becuase I was scared:o Anywho, I wanted to try to have my reactant be anhydrous ethanol. Added a few ml at a time and the solution turned yellowish green and fizzed very very little. Neglible heating was observed. After I had added al the ethanol I wanted(about 100ml) a brisk but not violent effervescence set in after a delay. All at once Cyanuric acid precipitated and the solution is clear with a nose burning fruity smell(acetylaldehyde?). I did not notice any chlorine or other simular fumes other than the pungent aldehyde sting in my nostrils. This really could be acetyladehyde on the cheap but hot do you separate aldehydes from ketones? I know ammonia bonds to aldehydes would this be the secret? I cannot use sodium bisulfite solution.


Closing comments: No chlorine was evolved at least where I could detect. The initial fizzing might have been the start of the reaction and enough hydrochloric acid byproduct had formed to force the reaction to completion. (Where is my pH paper?)

But I speicifically refrained from adding acid because I did not want chloroacetone and I did not get it. What is going here?




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[*] posted on 12-5-2006 at 16:32
Time to play at last


Well, I finally had the chance to try the copper sulfate and sodium dichloroisocyanurate. I noticed what seems like a rather complex change. When I first added the Dichlor no change in color was notice but a transcient oily layer formed on top and the solution assumed a couldy blue color that momentarily phased through blue green color VERY simular to cupric chloride not complexed with HCl. After the quite brief bluegreen phase, a series of increasingly darker blues finally turn purple then bright purple. I definately think there is complex formation here but what complex I do not know. I added considerable dichlor and the filtrate is a blue violet color. My Advanced Inorganic Chemistry book speaks of violet TRIVALENT copper complexes!! COuld this be. Maybe we are making sodium dichloroisocyanuro(III)cuprate.



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[*] posted on 10-7-2006 at 04:08


An Insight of the Reactions of Amines with Trichloroisocyanuric Acid
L De Luca, G Giacomelli
Synlett, pg.2180-2184,2004

Abstract:
The reaction between amines or a-aminoacids with trichloroisocyanuric acid is studied under various conditions: N,Ndichloroamines, nitriles and ketones can be obtained from primary amines, while free aminoacids undergo oxidative decarboxylation to the corresponding nitrile of one less carbon atom.

Key words: dichloroamines, trichloroisocyanuric acid, nitriles,
aminoacids

Attachment: An Insight of the Reactions of Amines with Trichloroisocyanuric Acid.pdf (160kB)
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[*] posted on 10-7-2006 at 18:39


Thanks solo! For once a TCCA reaction without exotic solvents/catalysts!:D I especially like the nitrile route.



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[*] posted on 12-7-2006 at 20:38


Quote:
Originally posted by woelen

I wonder, what interesting things can be done with the cyanuric acid, which remains. It is so easy to purify this, it seems like a waste to simply throw this away.

As you can read, I already did some experimenting with these chemicals (which I now have just a few days), but I would like to have suggestions from other members. Any ideas, but also any interesting facts about these chemicals are very welcome. I really think that these chems are interesting enough to justify a thread, devoted to them.


Definitely don't throw away the cyanuric acid as worthless .
In fact you can likely buy the cyanuric acid alone at the same suppliers , as it is used as a supplemental stabilizer .

The cyanuric acid remaining may be cheaply converted to
cyanurate salts of various metals , which can be thermally decomposed at red heat to give the metal cyanamide .

I have prepared the cyanurate of calcium for this purpose
but have not yet converted it to the cyanamide by thermal
decomposition .

One of the other threads you mentioned gives the details
of preparation for the calcium cyanurate , by neutralizing a hot suspension or solution of cyanuric acid with 2 equivalents
of NaOH to form the disodium cyanurate , and then running in CaCl2 solution to form Ca cyanurate which precipitates from the supernatant salt solution , via a simple double decomposition reaction ( metathesis ) . Magnesium and Zinc
cyanurates should result from the same scheme by use of their commonly available sulfates substituted for CaCl2 ,
although I have not yet tested these or other metallic salts .

Here is the link for the thread where this was described .
From previous experiments , the metal cyanurate would seem to be the most promising candidate as a direct precursor for the associated cyanamide , as no other product
should result from its pyrolytic decomposition , and this should be a convenient lab method for the synthesis of
pure cyanamides .

https://sciencemadness.org/talk/viewthread.php?tid=2762#pid5...
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[*] posted on 28-8-2006 at 13:20


Quote:
Originally posted by woelen
This indeed is a very nice experiment. With TCCA, however, the result is disappointing. I can hardly see any red glow. With Na-DCCA, the result, however, is lovely. That is the best chemiluminiscence experiment I've ever seen, besides the well-known luminol experiment. Probably Na-DCCA works better, because it dissolves easily in water, while TCCA is almost insoluble. Na-DCCA also still has over 60% available chlorine.
I also did this experiment by adding solid Ca(ClO)2 to 30% H2O2. This also gives a nice result, but with Na-DCCA the reaction is longer lasting. You have a nice glow for a longer time, while with Ca(ClO)2 there is a red glow for a fraction of a second, accompanied with a very violent reaction.


What procedure did you use for the chemiluminesence with Na-DCCA and peroxide? I tried just combining the two, which resulted only in lots of gas formation.
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[*] posted on 28-8-2006 at 13:46


Vetry simple. Take 2 or 3 ml of 30% H2O2 and add a spatula full of Na-DCCA. Do this, while the room is darkened. You'll see a lovely red glow. Try to avoid breathing the gas mix, formed in the reaction, it is quite pungent.



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[*] posted on 28-9-2007 at 15:49


what about C3HCl2N3O3 (sodium dichloro-s-triazinetrione hydrated)

is that good for anything?? 99% pure with 55.5-57% available chlorine




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[*] posted on 28-9-2007 at 16:02


it is moderately soluble in water. I mix it with lye and use as drain cleaner. Also, you should mix a solution of dichlo with a solution of Copper sulfate and tell me what happens. :D



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[*] posted on 29-9-2007 at 08:27


ok well i am wanting to try the a TCCA reaction. I have noticed that different acids are used to start the reaction HCl, H2SO4, what other options are there?

[Edited on 29-9-2007 by Slimz]




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[*] posted on 29-9-2007 at 09:09


HCl gets oxidized by TCCA to form chlorine gas. Try not to poison yourself.
About your question I can not answer since you did not explain what reaction and what role the acid is supposed to have.
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[*] posted on 29-9-2007 at 11:58


oops sorry.. i was referring to water/acetone and adding TCCA. then adding a few drops of acid to kick off the reaction.

would acetic acid work?

[Edited on 29-9-2007 by Slimz]




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[*] posted on 30-9-2007 at 01:22


Look what Nicodem said:

"HCl gets oxidized by TCCA to form chlorine gas."

If you'd read that you would see that acetic acid can't be used instead.
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[*] posted on 30-9-2007 at 06:11


Right.. well i need to obtain some sulfuric or hydrochloric acid then... knew i would need them eventually... The HW stores around dont sell muriatic acid any more.. The sell a crapy acetic acid substitute.

----edit----

Good ol nitrate test kit bottle #1 (for a fish tank ) is 41% hydrochloric acid

[Edited on 30-9-2007 by Slimz]




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[*] posted on 30-9-2007 at 10:47


Any acid that is able to catalyze the enolization of ketones will do as long as it does not get oxidized by TCCA. This reduces the choice to H2SO4, some Lewis acids and strong organic acids like CF3COOH, MeSO3H, TsOH etc. Like it was already said HCl gets oxidized by TCCA and so formed chlorine is a most unselective chlorinating reagent for acetone. High selectivity is what you want when the only practical separation technique is fractionating the product. And fractionating an extremely lachrymatory substance is no fun! Even without a column is the plain distillation a nightmare enough.

Whatever you do, do not do stupid things like "adding a few drops of acid" to a solution of TCCA in acetone!

UTFSE to see what kind of terrible things can happen if you do not follow the French patent word by word.
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[*] posted on 1-10-2007 at 12:02


ok so will this acid suit my needs?
http://www.onlinesciencemall.com/Shop/Control/Product/fp/SFV...




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[*] posted on 1-10-2007 at 13:28


This acid is perfectly suitable. The price, however, isn't. Over $5 for just 30 ml of acid? Try to find another source. Unfortunately I cannot be of any help here, because I live in the EU, but inside the USA there must be sources of sulphuric acid, which ship to private persons and sell acid in liter quantities.



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[*] posted on 2-10-2007 at 03:29


Isn't sulfuric acid available in hardware stores in the USA or something changed due to some new regulations? In EU anyone can buy diluted H2SO4 as used to refill car batteries. There are even sources for the concentrated one (though concentrating it to >90% is not particularly difficult). I don't understand why one would want to buy it from chemical resellers, especially considering their prices.
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[*] posted on 2-10-2007 at 05:39


Its not available in hardware stores in Australia, but it IS available in places like "battery world". And thanks to its use as an electrolyte it is quite pure too. It always seems to turn pale yellow when I concentrate it, but I cant rule out dust as the cause. Anyhow, even slightly tainted it is still good enough for most purposes, only when I am unsure do I fall back on the expensive lab quality stuff.
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[*] posted on 2-10-2007 at 06:04


ill check the battery warehouse and see if they sell it, but i have heard that most of the battery stuff is industrial waste. I would much rater have a pure sample than a galloon of crap.

THis looks ok
http://www.grainger.com/Grainger/categories/electrical/batte...

what do ya think?

[Edited on 2-10-2007 by Slimz]




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