Sciencemadness Discussion Board

TCCA, Na-DCCA and cyanuric acid

woelen - 24-4-2006 at 00:00

Last week I purchased 1 kilo of almost 100% pure trichloroisocyanuric acid (TCCA) and 1 kilo of almost 100% pure sodium dichlorocyanuric acid (Na-DCCA, a.k.a. sodium dichloroisocyanurate). Both of these are used as "stabilized chlorine" for outdoor swimming pools. In the Netherlands, both chemicals can be purchased in the pure state without any additives like NaHCO3 or Na2CO3 as "Pool power Choc" (Na-DCCA, granulate) and "Pool power Mini" (TCCA, tablets 20 grams). Probably this "Pool power" brand also is available in other countries.

These chemicals are quite interesting and can be used for generating chlorine gas in large quantities at fairly low price (between EUR 15 and EUR 20 per kilo, when purchased in 1 kilo quantities, much cheaper per kilo, when purchased in packs of 10 kilo).

Here, on sciencemadness, there is info about these chemicals, scattered over many different threads, but I think that these are so interesting that a special thread could be devoted to them.

I intend to use them for making chlorine gas, but a lot of other interesting things can be done with them. When I add hydrochloric acid to these chemicals, then Cl2 is produced, but a sludge of white cyanuric acid remains. When the sludge is boiled and all chlorine is driven off, then one can slowly add small quantities of distilled water to the hot liquid, until all solid cyanuric acid dissolves. On cooling down, lots of featherlike crystals of cyanuric acid are formed, which are easily dried and are absolutely free of smell of chlorine or swimmingpool-water.

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.

Na-DCCA dissolves in water very well, TCCA does not dissolve in water (or only very slightly). TCCA, however, dissolves in acetone very well and on evaporation, a mint-like smell remains! Is this due to formation of chlorbutol?

TCCA also does not dissolve in a solution of NaOH. It gives a white slurry. This surprises me. I expected to get a clear solution, with NaOCl and trisodium cyanurate in solution.

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.

solo - 24-4-2006 at 03:45

The TCICA Test for Distinguishing Primary and Secondary Alcohols
Gene A. Hiegel and Afshin K. Chaharmohal
J. Chem. Educ. 1997 74 423.

Abstract
Simple primary and secondary alcohols can easily be distinguished by their rate of oxidation with trichloroisocyanuric acid (TCICA).1,2 The TCICA test is conducted by adding the unknown to a solution of TCICA in acetonitrile3 containing hydrochloric acid4 and measuring the time for a precipitate5 to form. Primary alcohols react slowly and secondary alcohols react rapidly. To generate comparison data for primary and secondary alcohols, tests should be first carried out using ethanol and 2-propanol.6

TCICA Test Procedure

To a small test tube add 0.5 mL of the TCICA solution in acetonitrile (30 mg/mL), one drop of 1 N HCl, and one drop of the sample. After noting the time, flick the test tube with your finger several times to mix the contents. Watch the test tube until a precipitate forms and record the elapsed time. The elapsed time could range from a few seconds to a few minutes.

In order to destroy any unreacted TCICA before disposal in the appropriate waste container, a few crystals of sodium hydrogen sulfite (NaHSO3) and a few drops of water should be added to the test tube, and the mixture should be allowed to react for a few minutes.7

Caution: TCICA solution is a bleach and a strong oxidizing agent and will discolor clothes. Do not get the solution on your clothes, your skin, or any lab surfaces. Spills can be cleaned up with sodium hydroxide sulfite solution.



Notes

1. Trichloroisocyanuric acid can be used to prepare ketones from secondary alcohols; see Hiegel, G. A.; Nalbandy, M. Synth. Commun. 1992, 22, 1589.

2. Other methods for distinguishing alcohols include the Lucas testsee Roberts, R. M.; Gilbert, J. C.; Martin, S. F. Experimental Organic Chemistry; Saunders: New York, 1994; p 693; and 1H NMR see McGreer, D. E.; Mocek, M. M. J. Chem. Educ. 1963, 40, 358.

3. As other strong oxidizing agents, TCICA should be added to the solvent, not the solvent to the TCICA. Solutions of TCICA in acetonitrile are stable for years when stored in a brown bottle. TCICA is used as a swimming pool disinfectant and is widely available.

4. The reaction will proceed without the HCl, but the reaction times are more reliable when it is used.

5. The precipitate is cyanuric acid, a solid used to inhibit light-induced destruction of chlorine in swimming pools.

6. We have used this test in organic lab classes for several years, and it gives reliable results in the hands of students. Students are not provided with comparison reactions times; therefore, each student runs the ethanol and 2-propanol tests to get this information. Primary alcohols take about 7-30 min to give a precipitate and secondary alcohols take about 0.1-1.2 min. Tertiary alcohols, which cannot be oxidized directly, take 3 or more hours before a precipitate begins to form; presumably this occurs after dehydration to an alkene.

7. Iodide-starch test paper wetted with water can be used to test for the presence of oxidizing power before the solution is placed in the waste container

Attachment: The TCICA test for distinguishing primary and secondary alcohols.pdf (86kB)
This file has been downloaded 6218 times


Chris The Great - 24-4-2006 at 19:56

I tried adding TCCA to NaOH, or NaHCO3, and just got lots of clear, irritating and stinky gas. No clue as to what the reaction was. Anyone else have a clue? I was just expecting the same as woelen.

Interesting test there solo, easy except for needing acetonitrile.

praseodym - 24-4-2006 at 20:37

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. :cool:

woelen - 24-4-2006 at 23:05

Quote:
Originally posted by Chris The Great
I tried adding TCCA to NaOH, or NaHCO3, and just got lots of clear, irritating and stinky gas. No clue as to what the reaction was. Anyone else have a clue? I was just expecting the same as woelen.

Is your TCCA pure? I did the test another time and checked the presence of gas.
I only get a white slurry, no gas at all. The liquid alsmost remains odourless, just a faint smell of TCCA (typical swimming pool odour).
What happens if you add dilute sulphuric acid to your TCCA? No gas should be formed and the solid should not dissolve. If a colorless gas is formed, then the TCCA is not pure and contains "enhancers".

woelen - 25-4-2006 at 13:54

I did another set of experiments and did much more careful observations:

1) Add solid TCCA to a concentrated solution of NaOH: The solid breaks down fairly quickly. The liquid becomes turbid and light green. A very small amount of gas is formed, but not very much. Enough, however, to be observed easily. To my opinion, the gas is odourless. I could smell some typical swimmingpool odour, albeit weak, but this odour also is produced by the liquid, when there are no bubbles.
2) Add solid Na-DCCA to a concentrated solution of NaOH: Almost the same kind of observation as with solid TCCA.
3) Add a solution of Na-DCCA in water to a warm concentrated solution of NaOH: When the two liquids are mixed, then the liquid becomes very light green instead of colorless, also tiny bubbles are formed, but this is very minor.

I explain the observations as follows. This is not established chemistry, it is my own theory.

Both TCCA and Na-DCCA react with NaOH to form hypochlorite and sodium cyanurate. The latter is soluble in water, but in very concentrated NaOH it is less soluble. The green color is due to the formation of the hypochlorite. Hypochlorite ion is very pale green/yellow in solution, due to an equilibrium with the very weak acid HOCl, which is light yellow/green.

with TCCA: 6OH(-) + C3N3O3Cl3 ---> C3N3O3(3-) + 3ClO(-) + 3H2O
with Na-DCCA: 4OH(-) + C3N3O3Cl2(-) --> C3N3O3(3-) + 2ClO(-) + 2H2O

The Na(+) ions are just spectator ions.

At high Na(+) concentration, however, the C3N3O3(3-) ion (cyanurate) gives a precipitate of Na3C3N3O3.

The very weak evolution of gas probably is due to decomposition of ClO(-)/HOCl. At very high concentrated (e.g. near the solid TCCA or Na-DCCA in conc. NaOH-solution) a small part decomposes. Also, when the liquid is warm, some of the hypochlorite may decompose. Hypochlorite in solution is quite unstable, when its concentration goes beyond 15 ... 20 % active chlorine.

Polverone - 25-4-2006 at 18:38

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

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.

Cyanuric acid has been very useful to me in preparing cynates/cyanurates and subsequently cyanides by high-temperature reduction, but that is less useful if you have a ready commercial supply of cyanides.

I believe that cynauric acid can be used as an in-situ source of cyanic acid for ester formation in high-BP alcohols. It will certainly disappear after boiling in glycerol, presumably due to ester formation.

TCCA forms a lovely purple copper salt. I've never seen any other purple copper compound.

TCCA will rapidly vanish, bubbling and emitting visible vapors, in a solution of aqueous ammonia. If it weren't for the lachrymatory vapor or gas emitted it would be a neat stage trick, almost like dry ice but shorter-lasting.

Organikum - 25-4-2006 at 18:59

TCCA may be named as you wish "green" or watever. It will make you cry.

Trust me on this.

/ORG

woelen - 25-4-2006 at 22:34

Quote:
Originally posted by Polverone
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.

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.


Quote:
Cyanuric acid has been very useful to me in preparing cynates/cyanurates and subsequently cyanides by high-temperature reduction, but that is less useful if you have a ready commercial supply of cyanides.

I believe that cynauric acid can be used as an in-situ source of cyanic acid for ester formation in high-BP alcohols. It will certainly disappear after boiling in glycerol, presumably due to ester formation.

TCCA forms a lovely purple copper salt. I've never seen any other purple copper compound.

TCCA will rapidly vanish, bubbling and emitting visible vapors, in a solution of aqueous ammonia. If it weren't for the lachrymatory vapor or gas emitted it would be a neat stage trick, almost like dry ice but shorter-lasting.

A lot of interesting suggestions. Like the chemiluminiscence experiment, they sound interesting. I'll try them. Especially the purple copper salt sounds interesting. Is this simply prepared by adding a solution of copper sulfate to TCCA? Is it possible to isolate this salt? If you have a "recipe" for this, that would be nice.

In the meantime I also found a source for pure cyanuric acid (outdoor swimming pool stabilizer for hypochlorite based chlorine) for EUR 15 or so per kilo. Can this easily be depolymerized, making cyanic acid and cyanates? I also was thinking of making cyanogen chloride. Could a reaction of HCl (dry, made by adding NaCl to H2SO4) and cyanuric acid give cyanuric chloride (not to be confused with TCCA, it lacks the O's), which in turn can be converted to cyanogen chloride. This would open up a source of cyanide, without the need to use red heat on a mix of carbon and cyanates or cyanurates.

Polverone - 26-4-2006 at 00:06

Quote:
A lot of interesting suggestions. Like the chemiluminiscence experiment, they sound interesting. I'll try them. Especially the purple copper salt sounds interesting. Is this simply prepared by adding a solution of copper sulfate to TCCA? Is it possible to isolate this salt? If you have a "recipe" for this, that would be nice.

It actually seems to work better with the sodium salt rather than the pure acid. This may be due to pH effects or to poor aqueous solubility of TCCA. When I observed it before it was incidental to what I was doing so I didn't observe too closely. Just now I have tried again and I think I can say how: grind together TCCA and CuSO4 in a mortar, add a bit of water, and grind in sodium carbonate. The mixture will turn purple. This isn't how I originally observed it, and I think the color was stronger when I saw it before, but as I said I was not focusing on the purple compound at the time. Perhaps I should have paid more attention to it because whatever I was doing at the time did not pan out, while the memory of the purple compound stuck with me.

While trying to remember how to get the color just now, I found another interesting reaction. Grind together NaOH and CuSO4 with a bit of water so you have a strongly alkaline solution containing freshly precipitated CuO. Grind in TCCA and the black color changes to a rich reddish brown. Dilute with water and the mixture bubbles, reverting to black. I don't know if the decomposition on dilution was strictly from dilution; it could also be from the heat of NaOH dilution.

Quote:
In the meantime I also found a source for pure cyanuric acid (outdoor swimming pool stabilizer for hypochlorite based chlorine) for EUR 15 or so per kilo. Can this easily be depolymerized, making cyanic acid and cyanates? I also was thinking of making cyanogen chloride. Could a reaction of HCl (dry, made by adding NaCl to H2SO4) and cyanuric acid give cyanuric chloride (not to be confused with TCCA, it lacks the O's), which in turn can be converted to cyanogen chloride. This would open up a source of cyanide, without the need to use red heat on a mix of carbon and cyanates or cyanurates.

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.

garage chemist - 26-4-2006 at 07:39

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. :o

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]

woelen - 26-4-2006 at 12:01

@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.

----------------------------------------------------------------------------------------------------

@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 - 26-4-2006 at 12:52

The bottle says that it is 92% pure TCCA, this number is not the active chlorine content. However, it doesn't contain any carbonate fillers: no reaction with dilute H2SO4.
First I bought the wrong tablets, and they contained about 20% sodium carbonate (I downloaded their safety data sheet).
The tablets with carbonate fillers are sold as "fast- dissolving", while the ones made of nearly pure TCCA are sold as "slow- dissolving".

I am not sure if the reaction of TCCA with NaOH solution goes to completion, due to one reason: TCCA is industrially manufactured by adding chlorine to an aqueous solution of cyanuric acid in NaOH solution.
This means that the opposite reaction hypochlorite + cyanuric acid ---> TCCA also occurs readily.
There seems to be an equilibrium. However, the fact that some chloroform was produced in my experiment shows that some hypochlorite is present. However, the yield was low.

I have the suspicion that yields won't improve much by using the correct amount of NaOH (2,4g NaOH + 2,5g 92% TCCA).
Maybe only one chlorine atom of ther TCCA reacts, giving one mol of hypochlorite per mol of TCCA and Na-DCCA as byproduct.

I don't have Na-DCCA. It would be nice if you could try that out for me: 1 mol Na-DCCA, 4 mol NaOH, lots of water, some acetone, and see if there is chloroform being produced.

woelen - 26-4-2006 at 13:06

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]

woelen - 27-4-2006 at 11:03

Well, now it is time to do the experiment. I wonder, however, how you come to the ratio of 1 mole of Na-DCCA and 4 moles of NaOH. I find the following equation for the net reaction:

2CH3COCH3 + 3NaC3N3O3Cl2 + 8NaOH --> 2CH3COONa + 2CHCl3 + 3Na3C3N3O3 + 6H2O

For 1 gram of Na-DCCA, 0.4849 grams of NaOH and 0.176 grams of acetone are needed. If there is 100% yield, then I expect 0.3618 grams of CHCl3.

I'll perform the experiment in a long thin tube, if any CHCl3 settles, then I can see how many uL are formed.

Nevertheless, I'll also do the experiment with a 1 : 4 molar ratio of Na-DCCA and NaOH and see which works best. That means 0.72 grams of NaOH.

EDIT: results of experiments

First, I want to mention that all weights are +/- 0.01 gram.

I prepared three liquids:

A: 1 gram of Na-DCCA and 0.5 gram of NaOH, dissolved such that the total volume of the liquid is just below 20 ml.
B: 1 gram of Na-DCCA and 0.75 gram of NaOH, dissolved such that the total volume of the liquid is just below 20 ml.
C: 0.6 gram of aceton, mixed with 2.5 grams of water.

Liquids A and B both are clear and light green liquids. This is the maximum concentration I could achieve without the liquids being cloudy. I prepared them as follows:
Dissolve 1 gram of Na-DCCA in just over 10 ml of water. This results in an almost colorless and clear liquid. Do this two times, once for liquid A and once for liquid B.
Dissolve 0.5 gram of NaOH in a small volume of water and add all to one of the solutions of Na-DCCA. This results in formation of a white cloud in the liquid. Next, add water slowly, until the cloudiness just disappears. That works very well. The volume then is almost 20 ml. A similar result is obtained for liquid B, that also must be diluted to almost 20 ml before it becomes clear again.

Next, I did the experiments:

Add 1/3 of liquid C to liquid A and quickly stopper. The clear liquid becomes cloudy very quickly and considerable heat is produced.
Add another 1/3 of liquid C to liquid B. This has the same result as with liquid A.

Leave both cloudy liquids alone for 1 hour, while they are in a cool waterbath. After 1 hour, both liquids are still somewhat cloudy, but not as much as when the experiment started. There is a blob of chloroform at the bottom of the tube. There is no noticeably difference in yield for both experiments. My tube only allows measuring at 0.05 ml resolution and I could not read off easily. I estimate the volume of the blob to be appr. 0.05 ml. The yield is not very good. I expected around 0.2 ml of chloroform, so I only obtained 1/4 of what I expected.

As a final test, I checked whether all Na-DCCA is used up with 1/3 of liquid C. So, to one of the liquids I added the last 1/3 of liquid C. This does not result in heating up any more, so with 1/3 of liquid C I already had (slight) excess of acetone.

The blob definitely is chloroform, it had a nice sweet smell. Now, if we take into account the solubility of chloroform in water, then the yield is much better. Solubility of chloroform is 0.8 grams per 100 ml of water. I had a volume of appr. 20 ml, this means that 0.16 grams can be dissolved in that amount of water. If the blob indeed is around 0.05 ml, then there is another 0.075 grams of chloroform and the total yield then is 2.3 .. 2.4 grams. This is a yield of 2/3 of the theoretical amount, which seems fairly good.

Altogether, using Na-DCCA, chloroform can be obtained, but one has to put some effort in it to isolate it. Just collecting the liquid from the bottom results in too many losses. The chloroform, dissolved in the water, also must be distilled in order to have acceptable yield. The limiting factor is the fairly low solubility of the Na-DCCA/NaOH mix. Only 1 gram of Na-DCCA can be dissolved in almost 20 ml of water, to which quite some NaOH is added.

===============================================================

I also did some other experiments with Na-DCCA (polverone's copper experiments and making cyanuric acid). Some pictures are added here of results. The copper-salt looks really neat.

http://woelen.homescience.net/science/chem/exps/TCCA/index.h...

The color of the copper salt is really hard to capture. My digital camera has difficulties in reproducing the color precisely. What these pictures show is a good approximation, but you really have to see the precipitate yourself in order to appreciate its remarkable color. For copper this color is remarkable, because its normal color is in the greens, cyans and blues.

EDIT: fixed hyperlink, such that it works again.

[Edited on 12-8-12 by woelen]

garage chemist - 27-4-2006 at 13:42

Very good, woelen! Great results!
We have now reason to believe that all three chlorine atoms on TCCA can react with NaOH to produce hypochlorite and that this can undergo the haloform reaction.

The reaction mix must be distilled in order to obtain the chloroform, that's understood.
One should distill until only water comes over.
The chloroform must also be washed with some more water in order to remove the acetone, which should be present in a slight excess.

I also re-did my experiment with 2,5g TCCA and 2,4g NaOH, total volume of liquid 35ml (this becomes clear, no insoluble matter).
An excess of acetone was added under stirring, it became so hot that the produced chloroform started to boil, but I had my reflux condenser on it so no vapors escaped.
The blob of chloroform was distinctly larger this time.
I haven't measured yield, but it seems better than last time.
I don't have a means of mesauring small volumes, so the only real way to find out the yield is to run a large batch and distill the chloroform out of the mix.

garage chemist - 1-5-2006 at 11:35

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

Maybe I should have used somewhat less NaOH, as I now see that the reaction produces NaOH itself. Also NaOH can hydrolyse chloroform when the two come into prolonged contact, especially at elevated temperature.

Here the synthesis report:

The TCCA dissolved in the NaOH solution, however, it needed a few minutes of stirring. A small amount of foam was produced.
The acetone was diluted with two times its volume of water and put into a dropping funnel
The reaction was run in a three- neck 500ml round bottom flask with efficient reflux condenser.
The acetone solution was dropped in under violent stirring at a rate of about 3 drops per second.
It heated up rapidly, and soon started to boil. The rate of acetone addition was not decreased, but rather increased to bring the reaction to a finish as soon as possible (in less than 5 minutes), since heat causes hypochlorites to decompose.
As the acetone was added completely, the mix was left to cool down. Some chloroform was seen settling down.

The flask was rigged up for distillation, and slowly distilled under magnetic stirring. A chloroform/water- azeotrope came over at about 55°C, and as it stopped coming over the steam temperature actually went down until the residual solution started boiling. Then it quickly went to over 90°C, and only water came over. I stopped the distillation at this point.

To the raw chloroform in the receiver was added an equal amount of water and it was swirled to remove residual acetone.
The crude chloroform was extracted with a pipette and weighed.
It weighed 6g, which is exactly 50% of theory. Strange.
The distillation residue in the flask had absolutely no smell of chloroform, so nothing remained in there.
The chloroform was washed with conc. H2SO4 to remove the rest of acetone (the H2SO4 becomes dark red/brown when acetone is present), it was only discolored a slight bit.

I didn't post this in prepublication, because the yield is too low for this. The reaction will have to be optimised at first.

I can think of two explanations for the low yield:
first the high temperature in the synthesis could have caused half of the hypochlorite to decompose. Next time most of the reaction water will be added as ice, to provide internal cooling because with external cooling the reaction would take too much time to finish.
Second, I used too much NaOH. Next time I will use only 4 mol per mol of TCCA, to take into account the 2 moles being produced in the haloform reaction. This means that not all TCCA will dissolve when it is added to the NaOH solution.

A strange thing is that a white powder precipitated from the distillation residue, and it did so already while still boiling hot!
Cooling did NOT increase the amount of precipitate.
The liquid above it is also strongly alkaline, so it can't be cyanuric acid.
The residue was not present before distillation, when there was the chloroform.

[Edited on 1-5-2006 by garage chemist]

woelen - 1-5-2006 at 12:57

Did you isolate some of the white stuff, which was formed during distillation? This may me the key to explaining the low yield.

I'll see if I can reproduce that observation on a small scale. If that precipitate contains chlorine, then I can perfectly understand why the yield of CHCl3 is so low.

garage chemist - 1-5-2006 at 13:14

I still have the distillation residue in the flask, unaltered.

The precipitate only formed as the solution was being heated to boiling after the chloroform had been distilled off (although some of it may have formed during distillation of the chloroform, the mix was cluody the whole time).

I'll test the supernatant solution for chloride, with AgNO3 solution.
If there is really a lot of chloride in there, then it is clear that the chloroform has been hydrolysed by the excess of NaOH.

The reaction of chloroform with hot NaOH solution produces dichlorocarbene, an extremely reactive species which undergoes a large number of reactions. It may have reacted with the sodium cyanurate, forming some strange new
compound.

That makes good sense: the fact that the precipitate occured during distillation, and that the precipitate does not change solubility with temperature, e.g. is most likely insoluble, supports the theory that dichlorocarbene was produced from the excess NaOH and chloroform and reacted with the cyanurate.
I'll filter the precipitate and see what properties this material has. An unknown compound!

garage chemist - 2-5-2006 at 13:03

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.

woelen - 2-5-2006 at 13:46

Quote:
use the correct amount of NaOH: 4 mol per mol of TCCA

That is an interesting conclusion. I already found in my equation for Na-DCCA that I did not need 4 mol of NaOH per mol of Na-DCCA, but only 8 mol of NaOH per 3 mol of Na-DCCA. I derived the reaction equation at once, without the intermediate step of the hypochlorite. But, now we have explicit, what it the problem. The haloform reaction produces NaOH again, as you already mentioned.
Maybe you should even use a little bit less NaOH, than the computed amount. Just a little bit, to be sure that the liquid does not become alkaline.

Also, having an excess of acetone may be a bad idea. I noticed that acetone and TCCA react with each other, giving a refreshing and cooling smell (quite pleasant). I suspect that smell is due to formation of chlorobutanol. Chloroform and acetone also react with each other in the presence of an alkali (which acts as catalyst) as follows:

CHCl3 + OH(-) --> H2O + CCl3(-)
CH3COCH3 + CCl3(-) --> CH3C(O-)(CCl3)CH3

The latter ion reacts with water, giving CH3C(OH)(CCl3)CH3 and OH(-). This is isobutanol, with all three H-atoms of one of the methyl groups replaced by chlorine. A lovely very nice smelling compound, which makes you feel drowsy, but it is not what you want.

Next weekend, I am going to do some tests also with the correctly computed amounts of TCCA and NaOH. I already found out that TCCA does dissolve in NaOH quite well, but it does so very slowly. However, with patience and grinding it does dissolve. The only problem I observe is that part of it already decomposes while dissolving: formation of odourless and colorless gas, when the solid TCCA dissolves in NaOH-solution. Probably that gas is oxygen, due to decomposition of the ClO(-) at that high concentration.

garage chemist - 2-5-2006 at 14:22

Formation of chlorobutanol is very possible, in the alkaline liquid, and with excess of acetone. But my solution doesn't smell like it (I know the smell) and distillation would also get rid of it (it has a high boiling point).

With the stochiometry:

6 NaOH + C3N3O3Cl3 ----> 3 NaOCl + Na3C3N3O3 + 3 H2O

3 NaOCl + CH3COCH3 -----> CH3COONa + CHCl3 + 2 NaOH

We see that we can reduce the 6 mol NaOH to 4 mol, because 2 mol get produced later by the haloform reaction.
However, this means that the TCCA will not dissolve completely in the NaOH. I hope that this won't interfere with the reaction.

The chloroform will have to be fractionally distilled through a column after washing with conc. H2SO4, washing with water and drying with CaCl2 (this is what the Organikum recommends for purification of chloroform).

The whole thing would be quite labor- intensive, and only be worthwhile if done on a large scale (at least 50g TCCA).

Organikum - 3-5-2006 at 01:24

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

woelen - 3-5-2006 at 03:36

I did an experiment by adding some TCCA to acetone. It quickly dissolves and a clear and colorless liquid is obtained. There was, however, not a sign of any reaction. I let the acetone evaporate and a white solid was obtained again. This solid had a smell like mint/refreshing, the typical smell of chlorobutanol, but it was only very weak. Knowing that its smell is very strong, one can conclude that only a very small amount is converted. The solid also unmistakenly had the standard smell of TCCA (swimmingpool odour), so most of it is not converted at all. As a final test, I added dilute HCl to the white solid and it gives bubbles of chlorine and the turbid green/white liquid, typical of adding HCl to TCCA.

So, does the reaction between acetone and TCCA need a special initiation, or does it requitre heating to set off?

Organikum - 3-5-2006 at 07:00

The reaction of acetone with TCCA produces chloroacetones mono/di/ti depending on the amount of TCCA present. I posted a thread times ago about this method for the production of monochlorocatone.

The reaction is strongly exothermic.

Be careful!

A small amount of HCl helps initiating the reaction without having to apply heat whats a bad idea as it will end in a runaway. You have to cool it.

Is your TCCA scented?

/ORG

Nicodem - 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)
This file has been downloaded 1539 times


woelen - 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]

Organikum - 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

Nicodem - 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.

Organikum - 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.

Homebrew primary alcohol oxidation

chloric1 - 5-5-2006 at 14:19

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?

Time to play at last

chloric1 - 12-5-2006 at 16:32

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.

solo - 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)
This file has been downloaded 6155 times


chloric1 - 10-7-2006 at 18:39

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

Rosco Bodine - 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...

agent_entropy - 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.

woelen - 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.

Slimz - 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

chloric1 - 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

Slimz - 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]

Nicodem - 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.

Slimz - 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]

SecretSquirrel - 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.

Slimz - 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]

Nicodem - 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.

Slimz - 1-10-2007 at 12:02

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

woelen - 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.

Nicodem - 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.

Antwain - 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.

Slimz - 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]

not_important - 2-10-2007 at 08:50

Quote:
Originally posted by Slimz
ill check the battery warehouse and see if they sell it, but i have heard that most of the battery stuff is industrial waste...
[Edited on 2-10-2007 by Slimz]


Lead-acid batteries are fairly sensitive to impurities, particularly many metal ions, so the acid for them is fairly good grade but only 30 to 40 percent concentration.

Magnesium sulfate has been used as a user additive for lead acid batteries, it can squeeze a little extra life out of old batteries. There are other after-market additives, but they generally shorten the life of a new battery and/or reduce its performance.

Some manufactures specify that they use high purity acid
http://www.hbl.in/brochures%20pdf/VRLA%20leaflet-INDIA.pdf

but if you want a really good grade then you'll buy lab grade acid, or purify a technical grade yourself.

AJKOER - 22-7-2012 at 06:50

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


Here are a few scenarios as to possibly why the CHCl3 production yield is low. To start, I have an issue with the 1st reaction above as there is an implicit assumption on the stability/availability of HOCl given that the reaction actually proceeds in stages:

Initial Reaction:
NaOH + C3N3O3Cl3 ----> HOCl + NaC3N3O3Cl2

(support: see equation (1) at http://www.waterguardchem.com/tag/tcca/ which displays the reaction between TCCA and NaHCO3 forming NaC3N3O3Cl2 + CO2 + HOCl )

followed by:
NaOH + HOCl --> NaClO + H2O

but some of the hypochlorite may be moved closer to chlorate as a function of pH, temperature, concentration of HOCl and hypochlorite via the pH dependent reaction sequence (see discussion on 'Decomposition of Solutions' at "Handbook of Detergents: Production" by Uri Zoller, Paul Sosis, page 444, link http://books.google.com/books?id=dXn3aB1DKk4C&pg=PA444&a...):

when pH over 11:
2 NaOCl --> NaClO2 + NaCl and
NaOCl + NaClO2 --> NaClO3

or, much faster for pH 5 to 9, and slower but still faster than above for pH<5:
NaClO + 2 HOCl --> NaClO3 + 2 HCl

The HOCl itself may also disproportionate or decompose:

3 HOCl --> HClO3 + 2 HCl

2 HOCl --> O2 + 2 HCl

Also, HOCl can be attacked by the newly formed HCl:

HOCl + HCl <---> Cl2 + H2O

as is the TCCA and also the NaOH, all of which can impact the haloform reaction yield.

As support, note the comment from "Process for production of an alkali metal dichloroisocyanurate and trichloroisocyanuric acid", United States Patent 4395548:

"It has unexpectly been found that the inability to form sufficient hypochlorous acid to fully chlorinate the available dichloroisocyanurate to TCCA (a mol ratio of less than one) adversely affects each of these three factors. This low mol ratio can be caused by the decomposition of hypochlorous acid to form hydrochloric acid according to the following reactions: 2HOCl➝2HCl+O2 3HOCl➝HCL03 +2HCl, or by a reaction consuming hypochlorous acid, such as the formation of chloroamine: HOCl+NH3 ➝NH2 Cl+H2 O,
or by the consumption of available alkalinity as in the destruction of a triazine ring: 2C3 H3 O3 N3 +9Cl2 +18NaOH➝6CO2 +3N2 +18NaCl+12H2O"

See 3rd page, column 3, at link: http://www.google.com/patents?id=TNowAAAAEBAJ&printsec=a...


[Edited on 22-7-2012 by AJKOER]

AJKOER - 24-7-2012 at 14:46

I can across a even more likely candidate as to the poor CHCl3 production yield. Per a source (see "CYANURIC AND ISOCYANURIC ACIDS", Volume 8, page 202 at http://www.scribd.com/doc/30133467/Cyanuric-and-Iso-Cyan-Uri... ) to quote:

"5.2. Oxidation.
Although the triazine ring of cyanuric acid is stable to oxidizers such as peroxydisulfate, it can be cleaved by alkaline hypochlorite: 2 H3(NCO)3 + 9 ClO- --->3 N2 + 6 CO2 + 9 Cl- + 3 H2O (24). Chloroisocyanuratesare similarly decomposed by alkaline hypochlorite (8)."

Thus, Sodium dichloroisocyanurate formed by the initial reaction:

NaOH + C3N3O3Cl3 ----> HOCl + NaC3N3O3Cl2

may prove to be detrimental to any sequence NaClO formation.

Copper cyanurates

Bezaleel - 10-3-2013 at 19:06

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]

woelen - 11-3-2013 at 10:55

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.

AJKOER - 11-3-2013 at 17:59

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.


A couple of points:

1. TCCA + H2O forms, as previously noted, HOCl. In the presence of AgNO3, the reaction is given by (see http://chemistry.proteincrystallography.org/article176.html ):

3 HOCl + 3 AgNO3 = 3 HNO3 + 2 AgCl + AgClO3

I would ascribe this reaction as actually a two step reaction:

3 HOCl + 3 AgNO3 <--> 3 HNO3 + 3 AgClO

3 AgClO --warm-> 2 AgCl (s) + AgClO3 (this occurs rapidly)
-------------------------------------------------------
Net reaction, as noted above:

3 HOCl + 3 AgNO3 --> 3 HNO3 + 2 AgCl (s) + AgClO3

So my point is adding AgNO3 to a warm solution in the presence of Hypochlorous acid is not a valid test for chloride.
===========================================

2. HOCl dissolves in acetone and becomes relatively stable (limited disproportionation). This is a known procedures for extracting it from solution.

[Edited on 12-3-2013 by AJKOER]

Bezaleel - 12-3-2013 at 05:57

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.

Thanks, woelen.

Seeing the change in colour in the pictures above, it seems that merely the lines in the red part of the spectrum are disappearing on heating. This already gives some hint about what is happening.

The blue lines are characteristic of copper in aquaous environment. The red lines might be formed by interaction with nitrogen. Comparing the deep blue colour of copper nitrate in solution to the more sky blue colour of copper sulphate in solution, a similar process (also nitrogen involved) seems to happen there. But I'm guessing here. When I have found more information on this, I will post it here.

My guess is that the complex bound water molecules (6 for each copper atom) do not much change their arrangement when the complex is heated, but that the 4 cyanuric rings coordinated to the copper atom are then being rearranged in one way or another.

Refinery - 19-9-2014 at 05:41

My TCCA is mixed with sodium carbonate and I need to purify it. How it could be done safely? TCCA dissolves well in acetone and obviously there will be no reaction on room temp since there are oxidation reactions carried out with TCCA in acetone solvent.

I could dissolve TCCA in acetone, filter and vacuum distill off the acetone, but this is quite time consuming and still involves a risk of increased heat, because I dont have a cryogenic heat trap for sub-zero acetone.

But I need the TCCA dissolved in acetone anyway so I thought that I could measure the weight of acetone and TCCA, dissolve the mixture, filter off the residue and measure the mass of filtrate and solution to determine the approximate amount of TCCA recovered.

Refinery - 20-9-2014 at 10:16

In this video acetone and TCCA are boiled. So when in stable or pure form they do not react with each other? So can I mix TCCA with sodium carbonate to acetone, heat it, filter off the carbonate and distill off the acetone at water bath to make pure TCCA?

Metacelsus - 20-9-2014 at 11:54

Chlorinating agents such as TCCA can react with acetone to form chloroacetone. The reaction is autocatalytic due to HCl being a product, and the reaction being more favorable in acidic conditions.

If you do anything involving acetone and TCCA, test it on a small scale first.

Refinery - 20-9-2014 at 12:27

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]

AJKOER - 20-9-2014 at 16:03

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. :cool:

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.

20140714_101932_resized.jpg - 347kB

[Edited on 21-9-2014 by AJKOER]

j_sum1 - 18-12-2014 at 00:10

Reaction between TCCA and NaOH / KOH

Rather than start a new thread, I thought I would post my question here.
I have made the beautiful purple complex formed by reacting Na-DCCA with copper sufate. (Na2[Cu(C3N3O3Cl2)4)
I was interested in analogous complexes made by substituting the Cu with Co, Ni, Pb and whatever else I had on hand. I was also interested in substituting the Na for K or a group II metal ion.
To achieve this latter goal I attempted to make my own -DCCA compounds by reacting the appropriate hydroxide with TCCA.

I began by dissolving TCCA in NaOH solution in a 1:1 stoichiometric ratio -- hoping to achieve one dechlorination. TCCA dissolved and evolved some gas: no strong odour. Once dissolved I attempted to react with CuSO4. No purple precipitate but instead a light blue that looked like Cu(OH)2.
I tried the same with KOH. Dissolving took a long time and there was still some TCCA sediment. Again odourless bubbles were produced. Again no purple precipitate with copper sulfate -- just the light blue.
I concluded that I had not produced either Na-DCCA or K-DCCA, but I was unsure what reaction had occurred. My guess was that the bubbles were oxygen and that the solutions contained some unreacted hydroxide.

After an hour or so, both mixtures showed small flecks of bright purple on top of the Cu(OH)2 precipitate. I took this as a promising sign and possibly indicated that more time was needed to produce the -DCCA compounds. I decided to leave the NaOH/TCCA and KOH/TCCA solutions for a while and re-test later.
I came back five days later. There was a significant pool-chlorine type smell in the air but slightly different. Both the NaOH/TCCA and the KOH/TCCA were evolving bubbles which had a pungent odour that seemed not quite like chlorine. The sediment had fully disappeared from the KOH solution.
Adding CuSO4 gave no precipitate with either solution. This suggested that all of the hydroxide had reacted but also that there was none of the dichloroisocyanurate compound -- evidently I had overshot the reaction and was now getting some different products.

So, my question is, what are the reaction products from a hydroxide and TCCA? (0.1 mol/L at room temp). I am confident that excess NaOH would result in a triple dechlorination and leave cyanuric acid. I had assumed that by isng a 1:1 molar ratio the most stable product would be the dichloro species and that whatever equilibrium existed would favour that product. I wasn't expecting the gas evolution that I got. I wasn't expecting a second reaction destroying the Na-DCCA that I had produced.
Second question, what conditions might be required to effect a single dechlorination and maximise the -DCCA yield?

woelen - 18-12-2014 at 02:24

TCCA and NaOH do not simply give Na-DCCA and NaCl. This is not a matter of simple exchange of OH(-) and Cl(-). TCCA is in a more oxidized state than Na-DCCA (the difference is one equivalent of [O] between the two).

TCCA is C3N3O3Cl3.
NaDCCA is C3N3O3Cl2Na
You see the detail? In the latter, the Na has oxidation state +1, which is normal for Na. For TCCA the chlorines also have oxidation state +1, but this is a strongly oxidizing state of chlorine. Chlorine wants to go to -1, hence a difference of +2, which is the equivalent of one [O].


If you add NaOH to TCCA, then you get Na-DCCA and NaOCl. Formally, there is exchange of Na(+) and Cl(+). This is just formally, the actual mechanism is more complicated. The formal net reaction is

C3N3O3Cl3 + NaOH --> C3N3O3Cl2Na + HOCl

As you see, the liquid becomes acidic, by adding a base! This is a very very peculiar thing and is even more nicely demonstrated by adding ammonia to TCCA (be careful if you do that reaction in practice, it is EXTREMELY violent). You even get Cl2 and NCl3 under such conditions, compounds which only can be formed in acidic environments.

Hypochlorite and DCCA (or TCCA) react with each other, giving nitrogen gas and chlorinated derivatives of ammonia (chloramines). Finally, these two chemicals destroy each other and the material loses all its oxidizing power when sufficient NaOH is added.

Just to show you the reaction, take 1 gram of calcium hypochlorite and one gram of granular Na-DCCA (or 1 gram of granular TCCA). Mix the two solids and add one drop of water. Keep a distance from the mix, the reaction soon becomes extremely violent and the mix may even explode! DO NOT SCALE UP! Do this reaction outside.

As all of this shows, the chemistry of TCCA and friends is remarkable and very interesting from a theoretical point of view.

j_sum1 - 18-12-2014 at 03:06

Aaah. Makes sense. I hadn't considered the oxidation state of the chlorines.
That explains everything.

I was hoping to avoid chloramines. I guess I lucked out on that one. I will read up on these before I do too much more playing. I understand that they are a bit hazardous.

I can still play with Na-DCCA and try complexing with a few other transition metals. I recall you saying, woelen, that complexes of Pb, Co and something else have been documented. I must look up that post and the associated paper. I have a number of transition metal compounds to experiment with, however, most are oxides and carbonates since they come from a pottery supplier. I will need to convert to sulfates first. Time to make some more sulfuric acid.

Incidentally, I added some Na2CO3 to neutralise the solution before disposal. I have some extremely fluffy white crystals at the bottom of the beaker. I haven't seen any like them before.

Σldritch - 17-4-2016 at 07:35

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]

woelen - 18-4-2016 at 00:43

I purchased some pure cyanuric acid and dissolved this in a solution of NaOH (which is not easy at all, it dissolves very slowly). With this solution, however, I could not get the nice purple precipitate, when I add a solution of CuSO4. I only get a blue precipitate.

When I do the experiment with Na-DCCA to which some NaOH is added, then I do get the bright purple precipitate.

More investigation is necessary. Maybe lowering the pH of the solution of cyanuric acid does the trick, but my cyanuric acid only dissolves with great difficulty if the solution is not sufficiently alkaline.

j_sum1 - 18-4-2016 at 00:59

Hmmm.
This is one of my unfinished projects. I did get the purple precipitate after adding TCCA to KOH and then using copper sulfate. To say it was controlled would be a overstatement. I forget the exact details of my setup -- I will need to look at my notes.

One of the issues is that my TCCA is 30% borax. I don't know if it is making a difference but for this kind of experimentation I would prefer to be playing with pure compounds.

I have not had any success with cyanuric acid (which is pure.)

Quote:
The purple copper compound i suspect is actually copper (III)

No, the purple complex, (Na2[Cu(C3N3O3Cl2)4] is actually well understood and well documented. It is a well-known pool filter cloggant for people who randomly throw chemicals into their pool.

Σldritch - 18-4-2016 at 08:01

Ok, i have done the experiment again and this time i had more time to be more systematic, but first this:

Quote:
Quote:
The purple copper compound i suspect is actually copper (III)

No, the purple complex, (Na2[Cu(C3N3O3Cl2)4] is actually well understood and well documented. It is a well-known pool filter cloggant for people who randomly throw chemicals into their pool.


I may very well be wrong but my guess is that it is in a equlibrium between copper (II) and copper (III), first because of the unusual color and then because at the temprature of copper (III) oxide decomposition in water it turns blue wich, again i guess is due to copper (III) being less favored at higher tempratures. So basicly the color not only comes from copper (II) (The blue part of the color) but from a small amount of copper (III) wich gives the red color and together they make purple. I guess this is because the very oxidizing ligands copper has in this compound and that it is possible for a unpaired electron to delocalize over a big part of the molecule from what i can see. So maybe it releases a chloride ion and delocalizes an unpaired electron over itself. Also the result of the experiment below suggest were dealing with copper (III) even if its not a significant part of sodium copper Dichloroisocyanate.

EDIT: Forgot to mention that you might be able to test for cuprate (III) by comparing decomposition rate when using potassium hydroxide instead of sodium hydroxide due to the diffrent faster decompsition of potassiu cuprate (III).

Heres a helpful link: http://www.nrcresearchpress.com/doi/pdf/10.1139/v65-164

And my experiment in detail should be attatched.

[Edited on 18-4-2016 by Σldritch]

[Edited on 19-4-2016 by Σldritch]

Attachment: SodiumCopperDichloroisocyanurate+NaOH.odt (23kB)
This file has been downloaded 416 times


Fery - 23-9-2019 at 11:51

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.

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