TCCA, Na-DCCA and cyanuric acid

TCCA releases hypochlorous acid on contact with water. TCCA froms an explosive product with cyanuric acid + sodium hydroxide. Nitrogen trichloride, a highly explosive compound, may from if TCCA contacts ammonia, ammonia salts, urea, or similar nitrogen - containing compounds. This material is a powerful oxidizer and therefore, a potentially violent reaction with combustible materials may occur. Furthermore, an interesting property of TCCA is the chemiluminescence produced during their reaction with luminol in alkaline medium.

Quote:

Originally posted by praseodymFurthermore, an interesting property of TCCA is the chemiluminescence produced during their reaction with luminol in alkaline medium.

Metal cyanates (or at least sodium cyanate) can be made by heating an intimately powdered mixture of metal carbonate with cynauric acid, or so an old British patent told me. I never stopped heating early enough to try to isolate cyanates, as I was always interested in cyanides.

You can have some real fun with copper carbonate and cyanuric acid. Cyanic or cyanuric acid appears to form a bit more reddish purple compound with copper (oops, I guess I've actually seen maybe two sources of purple copper compounds, though quite similar if not indeed identical). You can see it if you trickle a bit of copper carbonate down the side of a test tube, then drop some cyanuric acid to the bottom and heat it over a flame. As the vapors come off they will react with copper carbonate clinging to the tube wall to form a purple compound like that near near the middle of the tube in the photograph.

In the case of the above photograph, I actually had excess copper carbonate powder sitting atop my initial load of cyanuric acid. I had thought it would be a good way to react more cyanic/cyanuric acid with copper carbonate before the vapors condensed again. But as I heated it, I saw no purple compounds in the bulk at the bottom of the tube, though there was evidently a reaction. Some acid vapors still escaped to react further up the tube as can be seen in the photograph. While I was continuing the strong heating, there was suddenly a little sizzling noise and a brief orange glow from the bottom of the tube in an area about the size of one of the original cyanuric acid grains. As I continued heating and rotating the tube there were a few more of these surprisingly vigorous reactions. Each one left behind metallic copper. I was only using the flame of an alcohol burner, and never expected copper carbonate/oxide to act in an almost pyrotechnic manner with cyanuric acid.

Nearer to the end of the tube there is a small dark blue patch, though the color may be difficult to see in the picture. It looks like ammonia-copper complex, and it's quite possible that cyanuric acid side reactions gave off ammonia.

Cyanuric acid goes to vapor easily but recondenses to solid easily as well. I don't know what would be necessary if you wanted to try to isolate liquid cyanic acid from it, even briefly.

The observation that TCCA dissolves in dilute NaOH solution made me repeat my experiment for the synthesis of chloroform from TCCA which didn't work when I last did it.

1,3g NaOH (0,03 mol + 0,1g excess) were dissolved in 30ml water and the solution cooled down with a cold water bath.
2,5g 92% TCCA (0,01 mol) were finely powdered with mortar and pestle (important!) and added to the NaOH solution under stirring.
Nearly everything dissolved, to yield a slightly green/yellow solution. It was decanted from undissolved material, cooled down again, poured into a 50ml ground- glass flask, a stirbar was added and it was stirred at medium speed.
A half pipette (ca. 1ml) of acetone was added and immediately a small vigreux column put on the flask, to serve as a reflux condenser so that no evaporation losses occured.

After about 30 seconds, the reaction suddenly kicked in, evident by it becoming suddenly cloudy and very warm.
The stirrer was switched off and slowly a blob of chloroform deposited at the bottom.

It works! One can make chloroform directly from TCCA, without using huge amounts of bleach!
I'm gonna run this reaction with twenty times those amounts, isolate the chloroform by distillation directly from the reaction mix and write an article in prepublication about it.

EDIT: I realize that I used only half of the necessary amount of NaOH. That may explain the low yield: it was far less than half a milliliter of chloroform.

[Edited on 26-4-2006 by garage chemist]

@polverone: I tried to make the purple compound and it works like a charm. What I did is dissolving the sodium salt Na-DCCA and add an excess amount of a solution of copper sulfate. When this is done, then a heavy precipitate with a very bright purple color is produced, which can easily be isolated. I rinsed the precipitate with water and now it is drying. The color really is very neat.

I also did the experiment of mixing basic copper carbonate with cyanuric acid. The cyanuric acid was made from Na-DCCA, to which excess dilute HCl was added and from which all chlorine was driven away by heating. On cooling down feather-like crystals were formed and the dried crystals I mixed with the copper carbonate. I, however, could not reproduce your nice results. I obtained a black powder on heating and at the colder parts of the test tube I obtained a white solid (evaporated cyanuric acid and resublimed again I suppose). I did not obtain all the reds, browns, purples and blues. Just black . The black stuff probably just is plain CuO. I also did not observe any of the 'pyrotechnic' reactions, although at a certain point I even heated the stuff with a propane torch.

I certainly will continue experimenting with copper / cyanuric acid / Na-DCCA or TCCA. If I obtain interesting results, then I'll certainly publish them on my website. A picture of the nice purple compound will be posted anyway, when it is dry.

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@garage chemist: That is a great job you did. I read the recent thread about TCCA and CHCl3 on versuchschemie.de and also the thread on sciencemadness with great interest, but it is even better that you actually succeeded in making CHCl3 as a separate liquid now from this compound. If you have optimized your preparation, then I certainly would like to see your article about that.

The TCCA you have, is it 92% TCCA or 92% active chlorine? I am inclined to think the latter, the number is too much of a coincidence: 100% TCCA has an active chlorine content of almost 92%.
My bottle tells that it has a minimum active chlorine content of 90%, so my TCCA probably is pure or almost pure. I did tests with H2SO4 and then I do not get any bubbles, and that is a good sign, so there are no carbonate fillers and the like.
My bottle of Na-DCCA tells me that it contains 100% sodium dichloroisocyanuric acid, and that on practical application at least 60% active chlorine is available in this compound. The theoretical active chlorine of Na-DCCA is 64%.

[Edited on 26-4-06 by woelen]

Garage chemist, right now, woelen is going in horizontal position (it is over 23:00 over here), but tomorrow evening I will try the experiment with Na-DCCA, NaOH and acetone and see what happens. I'll keep you updated with the result.

[Edited on 26-4-06 by woelen]

TCCA chlorinates acetone nicely, if I am not mistaken this proceeds up to tri-chlorination provided that enough TCCA is present.

The reaction is highly exothermic and the mono, di and trichlorinated acetone derivates are lachrymators (monochloro and dichloro are for sure).

So if I am not mistaken a way to do this might be to add acetone slowly to an excess of TCCA and after this has reacted to completeness to add NaOH to produce chloroform. One might even remove the precipitated cyanuric acid by letting it settle and decant the liquid - filtration is pretty futile - thus minimizing the amount NaOH needed.

I hope my memory didn´t fool me with the way the haloform works....

/ORG

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 . 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]

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.

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 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?

Time to play at last

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

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.

Quote: Originally posted by garage chemist

I have run the reaction with 24g NaOH (0,6 mol) and 25g 92% TCCA (0,1 mol) in 250ml water, and 7,2g acetone (0,124 mol). I figured that the reaction would go as follows: 6 NaOH + C3N3O3Cl3 ----> 3 NaOCl + Na3C3N3O3 + 3 H2O Then the haloform reaction: 3 NaOCl + CH3COCH3 -----> CH3COONa + CHCl3 + 2 NaOH

Copper cyanurates

Earlier in this thread, we spoke of the reactions between copper or copper ions and cyanuric acid or TCCA.
It is established in literature that copper will give a red-violet complex with sodium cyanurate, Na2[Cu(H2C3N3O3)4].6H2O. For an overview of cyanuric acid complexes, see the excellent review article by Seifer: Russ. Journ. Coord. Chem., 2002, vol 28 iss5, pp 301-324.

Personally, I would not say the complex is red-violet, but purple:


A peculiar thing happens when the substance is heated in water under stirring. The colour changes to a greyish blue, starting somewhere between 60 and 70°C. When the substance is cooled, the original purple compound is regained. The pH does not change when the colour changes, and lies constantly around 7. (Measured values were 6.7 to 6.8.)



Does anyone have any ideas what causes the change in colour?

[Edited on 11-3-2013 by Bezaleel]

Quote: Originally posted by garage chemist

I tested the liquid above the white stuff for chloride (added a few drops of it to some HNO3- acidified AgNO3 solution) and it was strongly positive (thick white clouds of precipitate). So either half of the hypochlorite disproportionated to chloride and chlorate (likely, as the solution was very hot during reaction) or half of the chloroform was hydrolyzed by the excess NaOH. Whatever it was, it is clear that I have to both a) keep the solution cold during haloform reaction, and add the acetone very slowly so that this is possible, and b) use the correct amount of NaOH: 4 mol per mol of TCCA.

Quote: Originally posted by woelen

It is quite special that the color returns to the original color on cooling down. It might be that there is a phase transition in the material (one solid form to another solid form) between 53 C and 85 C. I have done similar experiments with Ag2[HgI4] and Cu2[HgI4] where the color of the precipitate depends on the temperature and changes reversibly. A similar thing occurs with solid HgI2. All three experiments are on the following webpage: http://woelen.homescience.net/science/chem/exps/hgi2_thermoc... It might be that in your case there is a similar phenomenon, albeit less spectacular.

I figured it out that since the gunk contains almost 50% of sodium carbonate, it would neutralize any acids formed in situ, and since sodium carbonate is very slightly soluble in acetone, there will be always some. Quite the same way that adding .1% of NaOH into sodium cyanide will prevent it from decomposing to HCN in solutions.

I read that chloroacetone is stabilised by carbonates so it should be all way stabilising. Purity requirement is not very high - if it was, I would buy reagent grade TCCA, as long as it is decently pure. This TCCA is from my nearest mall so I bought it, but another supplier sells much more pure TCCA for pools which is about 90%. The rest of is an issue, though, it contains aluminium and copper sulfate, former being highly acidic and could catalyze a reaction, therefore I think it would be necessary to employ some sodium or calcium carbonate to neutralize it.

I gotta make some tests and write the results here so it could be useful for someone else looking for ways to purify TCCA.

My idea is to grind the tablets to fine and mix them with acetone and heat it up at water bath with reflux, and then let it cool a bit and filter everything that left over, and then distill off the acetone to get very pure TCCA. There should not be overheating issues since water bath can never go over 100C. I might consider adding a little amount of sodium carbonate to the filtered solution just in case it happens to go on it's own when distilling it. TCCA should dissolve up to 35g/100ml to acetone and maybe even more when heated.

Oh sorry I forgot to link the video:

https://www.youtube.com/watch?v=DDqIzxTGc78

[Edited on 20-9-2014 by Refinery]

[Edited on 20-9-2014 by Refinery]

Quote: Originally posted by Polverone

Quote: Originally posted by praseodymFurthermore, an interesting property of TCCA is the chemiluminescence produced during their reaction with luminol in alkaline medium. Another nice chemiluminescent reaction is the one between TCCA and hydrogen peroxide. It gives the red glow of singlet oxygen like the interaction between metal hypochlorites and peroxide, but TCCA offers more concentrated available chlorine (and thus more visible light) than commonly available hypochlorites. .....

Actually, I recently discovered the attached photo from my prior attempts on generating Singlet oxygen from dilute but cold NaOCl (8.25%) plus NaOH with a dilute 3% H2O2 drip. I didn' t think that the camera actually caught anything, but apparently, it did barely get something.



[Edited on 21-9-2014 by AJKOER]

The purple copper compound i suspect is actually copper (III) considering copper (III) oxide decomposes at about 75C and is coulored red. Considering the compound of NaDCCA and Copper Sulfate is purple it might be a mixture of copper (II) and copper (III).

I also did some experimenting with it by mixing it with an arbitary amount of hydrogen peroxide wich made the suspension of the copper compound turn white and bubble while the water turned a light blue.

The second experiment i did with it was mixing it with very concentrated sodium hydroxide wich gave very intresting results. Immidietly on addition of the purple copper compound it turn dark brown-red and the solution soon startes to bubble of what i think is oxygen gas (since it has only a slight smell of chlorine and metall). The solution was cold when starting the addition but increased slightly in temprature as it decomposed. I tried to filter half of the suspesion wich seemed to go trough the filter mostly but it did not seem to increase the rate of decomposition notably. The rest of the suspension i poured in portions into tap water wich gave an intresting effect: first it seemed to form a yellow orange solution wich within seconds turned to green to blue to a black suspension of what looked like copper (II) oxide.

Considering the red colour, the decomposition temprature, residue and the general instability i would make relativly un-educated guess that this is some form of copper (III); maybe percuprate seeing how it is only stable in extremely alkaline conditions.

Anyway this is my first post so please have mercy.

EDIT: https://en.wikipedia.org/wiki/Copper(III)_oxide

[Edited on 17-4-2016 by Σldritch]

While I was servicing basic military training we had Dikacid tablets in stock. If there was a war every soldier would obtain these tablets for dissinfection of water from unknown natural sources - to convert it to drinkable (microbiological aspect only, ineffective for neutralizing most of chemical poisons). These tablets were white when fresh, but as you can see, now they are already quite old. Every tablet contained 14 mg of sodium dichloroisocyanurate, now certainly less.
(picture 1, 2)
Recently I separated what to keep and what to trash and I rediscovered these tablets. So I started searching where to buy this substance now. Di-, tri-chloroisocyanuric acid and corresponding salts of analytical grade quality are expensive (about 100 EUR per 1 kg), but swimming pool grade quite cheap (10 EUR per 1 kg). I found these 2 in my country:
Wetter chlor start = sodium dichloroisocyanurate dihydrate 97% (small granules about 1 mm in diameter)
Blue pool mini tablets = trichloroisocyanuric acid 99% (tablets about 3 cm in diameter).

1st experiment:
(pictures 3, 4)
Dissolve 0,1 mol = 25,6 g of sodium dichloroisocynaurate dihydrate (26,4 g of Wetter chlor start) in 120 ml of water and filter into a 250 ml flask (remove undissolved impurities). It seems to be slightly endothermic process (just touched by my hand) - if do not dissolve completely, put into 40 C bath so it dissolves at the end.
Dissolve 0,025 mol = 6,24 g of Copper (II) sulfate pentahydrate in 30 ml of water in a beaker (if do not dissolve completely, put into 40 C warm bath) and add to the flask while swirling its content in a hand.
Purple precipitate is formed withing few seconds which settles very slowly in few minutes (the liquid above the precipitate was pale blue).
(pictures 5, 6)
Filter out the product, wash thrice with 10 ml of water, transfer it into a dish, let it to dry out at room temperature 20 C for 2 weeks (it is sticky and not willing to dry out fast).
CuSO4 + 4 NaC3N3O3Cl2 = Na2[Cu(C3N3O3Cl2)4] + Na2SO4
Yield of Na2[Cu(C3N3O3Cl2)4] 21,0 g.
(pictures 7, 8)

2nd experiment
(picture 9)
Dissolve 0,1 mol = 25,6 g of sodium dichloroisocynaurate dihydrate (26,4 g of Wetter chlor start) in 120 ml of water and filter into a 250 ml flask.
Dissolve 0,025 mol = 5,94 g of Nickel (II) chloride hexahydrate in 20 ml of water and add to the flask.
No precipitation occurred and reactants seemed to be vanished :-(
Dissolved 10 g of potassium chloride in 30 ml of water (endothermic, put into 40 C bath) and added to the flask while swirling its content in a hand.
A pale-green precipitate formed in the flask in few seconds :-)
(pictures 10, 11, 12, 13)
Filtered out the precipitate. Transferred the filtrate back into the flask. Washed the precipitate thrice with 10 ml of water and let it do dry out at room temperature 20 C for 1 week.
NiCl2 + 4 NaC3N3O3Cl2 + 2 KCl + 6 H2O = K2[Ni(C3N3O3Cl2)4].6H20 + 4 NaCl
Yield of K2[Ni(C3N3O3Cl2)4].6H2O 22,0 g.
(picture 14)
The yield could be slightly improved if I knew that Na2[Ni... is soluble. The KCl could be dissolved in in the solution of sodium dichloroisocynaurate so the final solution has 30 ml less volume.
The fun continues. Dual action, satisfaction. Dissolved 10 g of KCl in the 30 ml of washings, put the KCl solution into the flask with the filtrate and the flask into a fridge overnight (temperature +4 C) with a hope to get more product as the solution above the precipitate seemed to be still quite green (shifting the equilibrium to the right by increasing the concentration of K+ at the left side of the equation).
Few small crystals appeared at the bottom of the flask in the morning (approximately in a range of hundred milligrams) so I let the flask in the fridge for 3 days occasionally swirling the flask in my hand, there were slightly more crystals at the end. The attempt to increase the yield was not too much successful :-/ But every try counts :-)
The yield of crystals from the fridge was 1,1 g in slightly wet form. 1,0 g dry but already with signs changing color and 0,9 g after letting them in an open dish at room temperature 20 C and average humidity. Nice bright shiny pale green crystals changed color into almost gray.
This should be investigated further... maybe a hydrate with more molecules of water is formed at 4 C which decomposes to hexahydrate at 20 C?
(rest of pictures)

According US patent 3055889 it should be possible to obtain Li2[Cu(C3N3O3Cl2)4], Ca[Cu(C3N3O3Cl2)4], Ba[Cu(C3N3O3Cl2)4], K2[Cd(C3N3O3Cl2)4] also.