Sciencemadness Discussion Board - Nitrogen trichloride NCl3 production

Nitrogen trichloride NCl3 production

froot - 27-5-2004 at 11:23

I've seen discussions here involving this nasty stuff but a possibly stupid/interesting question still prods my curiosity.
I'm lead to believe that NCl3 can be made by bubbling Cl gas through NH4NO3 solution.
If one product is NCl3, could the other possibly be HNO3 under the correct conditions?
I'm wondering if there's a route using heat, a catalyst and a different solvent?

EbC: Edited title.

[Edited on 9-9-2006 by chemoleo]

Theoretic - 27-5-2004 at 11:44

Nooo... no heat, no catalyst, you're dealing with NCl3! I take to it you don't fully know the stuff... it's extremely heat/catalyst sensitive. The reaction you're proposing will run easily, in aqueous solution, and produce not simply HNO3, but a 3:1 (in moles) mixture of HCl + HNO3 - i.e., aqua regia...

S.C. Wack - 27-5-2004 at 12:05

Keep in mind that its discoverer, Dulong, lost three fingers and an eye from this. Davy, very aware of this, was knocked senseless by a small amount on his first cautious try, and was surprised by the damage done.

The_Davster - 27-5-2004 at 13:37

I have aparently once made some NCl3 acidently. After only being around it once for a very short period of time, I must say the sensitivity(even in solution) and how powerfull it is really surprises me.
Here is what happened http://www.sciencemadness.org/talk/viewthread.php?tid=2044

Proteios - 28-5-2004 at 22:31

Quote:

Originally posted by S.C. Wack
Keep in mind that its discoverer, Dulong, lost three fingers and an eye from this. Davy, very aware of this, was knocked senseless by a small amount on his first cautious try, and was surprised by the damage done.

The moral is..... think long and hard before doing anything that may be....irreversible.

Nitrogen Trichloride from Urea

mbrown3391 - 2-9-2006 at 06:36

Could anyone tell me the chemical reaction that occurs when urea is mixed with sodium hypchlorite (NOT in the presence of sodium hydroxide) that would yeild Nitrogen Trichloride? Also would this reaction have to be cooled, heated, or left at room temperture in order to yeild NCl3?

12AX7 - 2-9-2006 at 07:44

H2N-CO-NH2 + 7OCl- + H2O = CO2 + Cl- + 2NCl3 + 6OH- (I think that balances)

Since chlorine is removed by chlorination, charge is conserved by forming hydroxyls, so this will be driven by reducing pH, and should be stable at that (assuming it doesn't overreact and go high in pH, unchlorinating the product, but that doesn't really make sense). I mean, stable such as NCl3 is...

Don't know about thermodynamics.

Tim

mbrown3391 - 2-9-2006 at 09:30

If the reaction yeilds sodium hydroxide, what stops hydrazone from being formed instead of NCl3?

12AX7 - 2-9-2006 at 10:01

Hydrazone?? Would that be, erm, H2N-(N(+)=O)-NH2, i.e. the nitrogen analogue of acetone? But I've never heard of that so I assume you typoed hydrazine.

Should be that NH2- radicals form when the carbonyl is attacked (forming CO2), and they bond together to form hydrazine, which is further oxidized to N bits which chlorinate. Just a matter of carrying it farther.

Tim

mbrown3391 - 2-9-2006 at 10:13

Yes i did mean hydrazine. Would a large amount of household bleach be sufficient for this reaction, or would i have to chlorinate some sodium hydroxide? Also, i think household bleach contains traces of sodium hydroxide left over from production. Would this effect the reaction? Finally, at what tempurture should the reaction be kept?

mbrown3391 - 2-9-2006 at 10:40

Quote:

Originally posted by 12AX7
H2N-CO-NH2 + 7OCl- + H2O = CO2 + Cl- + 2NCl3 + 6OH- (I think that balances)

The oxygen doesn't balance actually. There are 9 oxygen atoms on the left and 8 on the right. Maybe it's this?:

H2N-CO-NH2 + 6OCl- + H2O = CO2 + 2NCl3 + 6OH-

12AX7 - 2-9-2006 at 11:44

Thanks, that hanging Cl didn't quite make sense.

Of course you can also add another H2O and form CO3(2-) + 2H+, killing two OH-, but it's still a pretty basic result (4 mol OH-, plus an ionized carbonate).

Tim

Marvin - 2-9-2006 at 13:08

Not that this isn't a wonderful game of scrabble you are having here....

mbrown,

Why do you think mixing sodium hypochlorite with urea will produce nitrogen trichloride?

mbrown3391 - 2-9-2006 at 13:25

Searching "urea" +"nitrogen trichloride" on google brings up countless pages that say mixing urea with sodium or calcium hypochlorite will form nitrogen trichloride. Heres one example:
http://www.jtbaker.com/msds/englishhtml/U4725.htm

Quote:

Incompatibilities:
Urea reacts with calcium hypochlorite or sodium hypochlorite to form the explosive nitrogen trichloride. It is incompatible with sodium nitrite, gallium perchlorate, strong oxidizing agents (permanganate, dichromate, nitrate, chlorine), phosphorus pentachloride, nitrosyl perchlorate, titanium tetrachloride and chromyl chloride.

Upon mixing small amounts of Clorox bleach and urea, i did observe the evolution of a gas. As 12AX7's chemical equation shows, CO2 gas would be evolved when urea and sodium hypochlorite are mixed. At the very least, a chemical reaction did occur. So, if it does not produce nitrogen trichloride, what does it produce?

Marvin - 2-9-2006 at 16:53

No, when you arrange the symbols to match the chemistry, that is a chemical equation. When you arrange letters for the purposes of getting things you recognise, that is scrabble.

An MSDS will list a lot more than quantitive reactions, it will list side products and rare events.

The actual products will depend a lot on circumstances, concentration, ratio of reactants. Simply thrown together in the 'expected' ratios for nitrogen trichloride, I would expect large amounts of carbon dioxide and nitrogen gas, some liberated chlorine with a possiblity of a non zero amount of NCl3. Assuming the mixture can get acidic enough. NCl3 forms only in slightly acid mediums and is removed by OH- or by reaction with ammonia, or urea.... Reliable means for making NCl3 are in another thread.

Classically the Urea and Hypochlorite reaction is the industrial synthesis of Hydrazine, a modification of the Rashig process.

[Edited on 3-9-2006 by Marvin]

JohnWW - 5-9-2006 at 13:54

Does anyone in New Zealand want some urea? I recently bought a box of it in an auction in small plastic bags, intending to use it as a fertilizer. and I have some surplus.

Nitrogen Trichloride through Electrolysis

mbrown3391 - 6-9-2006 at 16:54

I have read that it is possible to produce nitrogen trichloride through the electrolysis of an aqueous solution of NH4Cl. Can anybody explain the reaction which would occur during this process, if it is even possible?

hodges - 6-9-2006 at 17:32

See
http://www.sciencemadness.org/talk/viewthread.php?tid=2397

Hodges

hodges - 6-9-2006 at 17:37

Doesn't hurt to post this link as to what can happen when you accidentally form nitrogen trichloride (from hypochlorite and urea, no less):

http://www.chemaxx.com/explosion16b.htm

Hodges

mbrown3391 - 8-9-2006 at 15:20

I used followed your link, hodges, and i found this:

Quote:

"There is another method of NCl3 producing : from NaOCl , CH3COOH and NH4NO3(any ammonium salt)
The reaction is:
3NaOCl + 2CH3COOH + NH4NO3 = NaNO3 + 2CH3COONa + NCl3 + 3H2O
The reaction doesnt need any control - you just mix the solutions and wait 1-2hours , besides you dont need to work with gaseous chlourine A 5-7% NaOCl solution is used in Russia as a bleach
the method

make it outdoors because of extremely bad smell!!!)
reactives :

300ml of 5-7% NaOCl solution ,
50ml of 70% CH3COOH
solution of 10gNH4NO3 in 30ml of water

1)cool the solutions in refredgerator (about -10C)
2) pour 300ml NaOCl into a 0.5l plastic bottle(cola)
3) add 25ml of CH3COOH to NaOCl by small parts
4) add the ammonium nitrate solution to the bottle by small parts(extremely bad smell ! the solution must become yellow)
5) add 25ml CH3COOH to the bottle by small parts
6) leave the bottle(dont close it !) for 1-2 hours The solution must become colourless and a yellow-orange liquid must gather on the bottom (about 10g)
Then close the bottle and throw it somewhere"

I tried mixing small ammounts of the above chemicals and i did see a yellow precipitate. I then dripped a single drop of the supposed NCl3 onto a peice of paper and put a drop of turpentine on top of that. Nothing happened. I tried adding terpentine directly to the test tube. Still nothing happened.

Can anyone verify that the above chemical equation is accurate? If so, why can't I get the extremely senative NCl3 to detonate?

hodges - 8-9-2006 at 15:41

Mbrown, what did you use for your 70% CH3COOH?

You may be lucky it didn't detonate - I've heard stories of a tiny drop shattering containers and removing fingers when it detonates. I certainly wouldn't want my body anywhere near it when it detonates - which is why I placed oil on a long stick for my experiment, and even that a very small calculated amount based on the amount of current passed through during electrolysis. With larger quantities like it sounds like you were making I would not even want to be that close.

You may want to try making nitrogen tri-iodide instead. This is (relatively) much safer. There is plenty of information on nitrogen triiodide here and on the internet in general.

Hodges

mbrown3391 - 8-9-2006 at 15:52

I substituted the correct amount of 5% white vinegar for the 70% acetic acid. I thought it might not work, but since i did see a yellow precipitate i assumed that it.

Don't worry, i wasn't following those instructions exactly i used about 1 ml of bleach, 2 ml of vinegar and 1 ml of ammonium chloride.

And i have already made Nitrogen tri-iodide multiple times. It is just as interesting, but much more expensive as it requires iodine crystals.

hodges - 9-9-2006 at 18:21

I think the viengar is your problem - not nearly strong enough. You might want to try muriatic acid instead. There is nothing special about viengar in the reaction - I'm sure it was just chosen for expediency. You could use sulfuric acid as well but that would be even more dangerous because of all the heat it gives off when diluting.

Hodges

mbrown3391 - 9-9-2006 at 19:10

What would the chemical equation be, and is the HCl just a catylist or part of the reaction?

mbrown3391 - 9-9-2006 at 21:08

The case is closed on Nitrogen Trichloride production. There are many methods for creating it, most of which have been outlined in this topic, and for me all failed except for this one:

HCl is reacted with ammonia to produce NH4Cl. Additional HCl is added which (I assume) will later act as a catylist. A small amount of NaHCO3 is added to be sure the liqud contains excess HCl. The liquid is then placed in a plastic testube or beaker (NO GLASS!) and placed in an icebath. When the liquid is cooled well below room tempurature, 6% NaOCl is slowly dripped in. Immediately the liquid should become yellow and cloudy. Continue to add bleach until the liquid is completely opaque. Allow it to settle in the icebath. Within 5 minutes you will begin to see a distinct seperation between a cloudy layer and a transparent layer. The cloudy layer is NOT the finnished product! Within fifteen minutes you should begin to see tiny distinct droplets of a yellow oil on the bottom of the container. When the reaction is complete the solution will be completely clear with droplets of oil on the bottom.

I beleive the equation would be
-----------------------HCl
3NaOCl + NH4Cl = NCl3 + NaOH + HCl

NCl3 + NaOH + HCl = NCl3 + NaCl + H2O

If the reaction gets out of hand dump it into a cup of ammonium hydroxide, which destroys the NCl3

I beleive one could make NI3 by mixing NCl3 with KI, but i may be wrong.

NCl3 + 3KI = NI3 + 3 KCl

The advantage of this is that Potassium Iodide is much more readily available than Iodine crystals.

[Edited on 10-9-2006 by mbrown3391]

[Edited on 10-9-2006 by mbrown3391]

[Edited on 10-9-2006 by mbrown3391]

[Edited on 10-9-2006 by mbrown3391]

neutrino - 10-9-2006 at 07:08

Your equations don’t balance.

I’ve never seen this synthesis before, but It doesn’t look too reliable. Ammonium + bleach in a neutral solution producing NCl3? I just can’t picture that.

>The liquid is then placed in a plastic testube or beaker (NO GLASS!)

NCl3 explodes violently on contact with any organics. These instructions look like BS to me.

mbrown3391 - 10-9-2006 at 08:14

Sorry, i meant to type "a glass testube or flask (NO PLASTIC!)". Although this reaction would most likely not cause an explosion should it have been carried out in a plastic container. As the NCl3 is produced it is destroyed by the plastic so no amount can collect. I see the mistake in my equation, it should be

3NaOCl + NH4Cl = NCl3 + 3NaOH + HCl

NCl3 + 3NaOH + HCl = NCl3 + 2NaOH + H2O +NaCl

And the solution is not neutral, it's acidic.

I will post pictures soon.

mbrown3391 - 10-9-2006 at 09:20

As you can see, yellow droplets of NCl3 formed when I combined Ammonium Chloride and Sodium Hypochlorite in the precense of HCl.

hodges - 10-9-2006 at 12:05

Interesting pictures, I don't think I've seen NCl3 before. What happened to it after you took the pictures? I believe it is supposed to decompose on its own if it does not detonate first.

Hodges

woelen - 10-9-2006 at 12:47

Yes, this synth of mbrown3391 DOES work! I've done it myself many years ago. As I have NH4Cl, I simply dissolved some NH4Cl in dilute HCl (it does not dissolve well in conc. HCl) and then added NaClO (10%) drop by drop. I obtained a few drops of a yellow oil, with a very pungent, somewhat spicey smell (once you've smelled it, you'll recognize it the rest of your life). These drops are very much like the drops shown in the pictures above, so I think mbrown3391 did a good job making NCl3.

I just did the experiment and destroyed the stuff by slowly increasing the pH of the liquid. I was too scared to do real experiments with it

. At high pH it starts bubbling and disappears, I think it forms N2 and hypochlorite at higher pH.

Using NCl3 to make NI3 from KI seems like idiocy to me. If you really want to make NI3 from KI, then simply dissolve the KI in dilute HCl, and add excess dilute H2O2. Iodine is formed and quickly settles at the bottom. Rinse well with water, no need to dry. To this, add concentrated NH3 and you get your NI3.

mbrown3391 - 10-9-2006 at 12:49

After i took the pictures, i detonated it using the following setup. As you can see i only had a drop of it, so the explosion was not very big. It made a cracking noise, like NI3.

The explosion completely destroyed the cup, but did not damage the board.

Also i dont think the equation i posted previously is correct, as when i mixed the reactants i observed a lot of bubbling. Does anybody have any idea what the correct equation would be?

[Edited on 10-9-2006 by mbrown3391]

The_Davster - 10-9-2006 at 12:53

I too did this once a year or so ago, just made a miniscule drop though. Will try it again as soon as I can find my soft plastic pipettes and dig up the plastic glassware from an old chemistry set(what do you know, it DOES have a use).

mbrown3391 - 10-9-2006 at 12:56

On my first try, i immediately destroyed the NCl3 with NH4OH. This works pretty quickly. I am wondering how its discoverer lost three fingers and an eye from this stuff. How much did he make and what made it explode?

mbrown3391 - 10-9-2006 at 12:58

Rougue Chemist, you cant make this stuff in plastic, i tried it. as the NCl3 settles to the bottom, it reacts with the plastic and becomes something other than NCl3. I made this mistake the first time. Unfortunately glass is the only other option.

The_Davster - 10-9-2006 at 13:03

The stuff detonates in contact with most anything. Even a fingerprint on the glass will cause it to detonate(supposedly). Back then chemists worked with larger ammounts of chemicals, if you read the older preparatory books you will find all sorts of examples of quite dangerous chemicals being made large scale. Like the preparation of 10g pure Mn2O7

EDIT: Really? reacts with plastic, hmm, how unfortunate.

[Edited on 10-9-2006 by rogue chemist]

mbrown3391 - 10-9-2006 at 13:06

So what is the largest amount you would prepare and detonate at one time?

EDIT: Yes. The first time i did it in a styrofoam cup (dont ask me why) and it ate right through the bottom. Then i tried it in a soda bottle. This time it didn't eat through but no drops formed, it just turned the bottom of the bottle yellow.

[Edited on 10-9-2006 by mbrown3391]

The_Davster - 10-9-2006 at 13:16

Well pretty much anything organic eats through polystyrene, hope I don't have to break out the teflon here.

As for how much one should make, that is up to you, taking into consideration how much damage you could do to yourself in the worst case scenario. In my book I probably would not try to go above 0.2 ml. After all, this is a novelty, not for practical use.
Even the old preparatory books don't make NCl3 pure, they make and keep it dissolved in carbon tetrachloride. When they take such precautions, you know it is serious stuff.

[Edited on 10-9-2006 by rogue chemist]

hodges - 10-9-2006 at 13:25

In my electrolysis experiment a while back I used a plastic container. That may explain the lack of any yield.

Hodges

mbrown3391 - 10-9-2006 at 13:36

Thats was definately the problem. Common sense tells you to use plastic so you dont get shrapnel if there was an accident, but in this case you have to use glass.

neutrino - 10-9-2006 at 18:17

I always assumed that you needed Cl2 to make NCl3. In woelen’s procedure, this is made in situ from the acid + chloride + hypochlorite.

I don’t see where the chlorine comes from in your procedure, though. Maybe it isn’t necessary after all?

mbrown3391 - 11-9-2006 at 11:07

I think this is actually the correct equation:

3NaOCl + NH4Cl + HCl = NCl3 + Cl2 + 3NaOH + H2O

Much of the free chlorine produced escapes but some will react in the following manner:
Cl2 + NaOH = NaOCl + HCl

Cl2 + H2O =HClO + HCl

2HCl + 2NaOH = 2NaCl + 2H2O

As you can see, the reaction ultimately produces NaOCl and HOCl along with the initial NCl3. Both would react with the remaining ammonium chloride to produce more NCl3. After one round of this cycle, the products are:

NCl3
NaOCl
HOCl
2NaCl
3H2O

This appears to allow the reaction to last much longer than it normally would. This is good for high yeild, but in a labratory experiment it is dangerous, making it difficult to judge how much sodium hypochlorite you should add initialy.

[Edited on 11-9-2006 by mbrown3391]

mbrown3391 - 11-9-2006 at 12:21

im trying to determine some exact ammounts here, but i dont know the percentage of NH4OH in Blue Ribbon Ammonia. Does anyone have any idea?

hodges - 11-9-2006 at 13:49

It seems that household ammonia solutions tend to be 2M to 3M concentration. I recall titrating some grocery store clear ammonia solution once and finding it was just over 2M.

Hodges

neutrino - 11-9-2006 at 14:24

I'm not sure where your equations come from. It's more like this:

First chlorine is formed in the acidic environment. Note that commercial bleach already contains more than enough chloride ions.
2H+ + Cl- + OCl- --> Cl2 + H2O

This is the step I’m not positive about, but it looks about right:
3 Cl2 + NH4+ --> NCl3 + 4H+ + 3Cl-

Overall,
3OCl- + NH4+ + 2H+ --> 3H2O + NCl3

That's why I question a neutral synthesis.

[Edited on 11-9-2006 by neutrino]

mbrown3391 - 11-9-2006 at 15:03

I know that a large amount of free chlorine is formed from this synthesis that is never part of the creation of NCl3, because it goes into the air.

Quote:

3OCl- + NH4+ + 2H+ --> 3H2O + NCl3

Your eqution does not include this.

EDIT: Unless maybe all the NCl3 is made in the few seconds it takes for the chlorine gas to rise out of the liquid? If this is the case, then yeild is severely lowered. Maybe adding bleach very slowly with a pipette into a flask stopped with a one-hole stopper would help. This way almost all of the chlorine could be absorbed.

[Edited on 12-9-2006 by mbrown3391]

woelen - 11-9-2006 at 22:31

Neutrino's equation is a net equation, and it seems correct to me. The intermediate is chlorine gas, but not all of it reacts. But that does not invalidate the net equation given. If a closed system were used, with excess NH4Cl and acid, then all hypochlorite would be used up in formation of NCl3.

There are, however, side reactions, such as formation of nitrogen gas, and possibly some hydrazine as well.

2bob - 13-9-2006 at 05:22

is it possible to make NaNO2 from ammonia?

I was thinking along the lines that if H2O2 + NH4OH = NH3NO2 + H20, then wouldn't NaOCO2 (Sodium percarbonate: in all the new bleaches) which breaks down to H202 + NaCO2 (in aq soln), break the bond between the NH3 & NO2 giving NaNO2 + NH3CO2 (which would precipitate, wouldn't it?) leaving pure NaNO2 in aqueous solution?

I am unsure of the chemistry here, however it appears possible from my somewhat limited experience, and possibly overly imaginative perspective?

12AX7 - 13-9-2006 at 06:51

NH4NO2 = N2 + 2H2O. (Oh, and subtle difference: it's aich-two-oh (H2O), not aich-twenty.

)

What is NH3NO2, NaOCO2 and NH3CO2?

I'm sure you meant NH4, Na2CO3.H2O2 (I forget how much H2O2) and (NH4)2CO3.

There is a true sodium percarbonate NaCO3 or so, and similarly perborate as well, but they are NOT the common adduct sold as "oxiclean", et al.

Ammonium carbonate is soluble; sodium bicarbonate is probably one of the least soluble alkali carbonates.

To oxidize something, you have to go through the intermediate states. You might get something like NH3 + 1.5 H2O2 = 3H2O + N, where N is a nitrogen radical or so (N(0), which readily combines with others to form diatomic nitrogen N2). The radical may combine with more oxidizer, for instance N + 2H2O2 = 2H2O + NO2, which disproportionates in solution 2NO2 + H2O = NO3- + NO2- + 2H+, which oxidizes further as NO2- is sensitive. But I doubt this reaction is particularly common in solution.

More often, ammonia is burned with oxygen on a hot catalyst, where such reactions are encouraged by the freedom of thermal motion and the action of the catalyst. The process is something like 2NH3 + 5O = 2NO + 3H2O, where NO is formed because NO2 disproportionates at this temperature (NO + O <--> NO2, which proceeds to the right at lower temperatures).

Tim

woelen - 13-9-2006 at 13:42

Forget about making NaNO2 from ammonia in a simple aqueous redox reaction. If ammonia is oxidized in aqueous solution (and it isn't by the peroxo/carbonate adduct you mention), then usually N2 is formed, sometimes hydrazine or hydroxylamine.

NaCO2 is something non-existent. You probably mean Na2CO3 (sodium carbonate). The sodium percarbonate you mention is a compound, which can be written as Na2CO3.xH2O2.yH2O (I don't know the precise values of x and y). When this is dissolved, you get free H2O2, sodium carbonate and water.

2bob - 14-9-2006 at 14:23

ta.

mbrown3391 - 14-9-2006 at 15:50

What does making NaNO2 from ammonia have to do with nitrogen trichloride?

Dornier 335A - 26-3-2014 at 07:07

I have tried synthesizing tiny amounts of NCl₃ a number of times now. Every time the liquid has turned cloudy yellow (like in mbrown3391's photos), but no drops have collected on the bottom. I have tried several processes including:
24% acetic acid mixed with 3% NaClO solution in 1:20, 1:10 and 1:1 volumetric ratios respectively, followed by addition of concentrated urea solution drop by drop. Temperature was approximately 15°C.
NH₄Cl dissolved in 24% acetic acid, followed by addition of 3% NaClO solution drop by drop. This was attempted at both 15°C and -5°C.
To make sure the acetic acid wasn't the problem, I switched to NaHSO₄ solution as another weak acid and performed the above procedure again, this time at -2°C.
Every time the yellow cloud forms when the solutions mix, but the solution turns clear, sometimes slightly yellow after some time, and almost immediately if I swirl it around carefully.
I'm not sure how to continue from this. Am I doing some obvious mistake?

TheChemiKid - 26-3-2014 at 07:12

I tried making NCl₃ once, I luckily did it in a fume hood with the hood down.
Long story short: My beaker blew up, and damaged the fume hood.

Dornier 335A - 26-3-2014 at 11:24

Then you obviously succeeded with making the NCl₃ at least. What procedure did you use?

The Volatile Chemist - 12-8-2015 at 07:55

I attempted the [hydrochloric acid] acidified ammonium chloride solution and hypochlorite path, as well as adding ammonium chloride solution to hydrochloric acid and adding manganese(IV) oxide to catalyze chlorine production, but neither have produced NCl₃. For the first method, I tried using many different concentrations of both acid, ammonia, and hypochlorite. Should I be using a lot more hypochlorite than the other two reagents? Or should a little NCl₃ form regardless?
For the second reaction, only copious amounts of chlorine were formed.
Are there any sure-fire methods that produce NCl₃?

electrolysis

franklyn - 12-8-2015 at 10:09

nitro-genes - 12-8-2015 at 17:39

TCCA and ammonia salt, try extremely small amounts only

[Edited on 13-8-2015 by nitro-genes]

blogfast25 - 12-8-2015 at 18:30

Quote: Originally posted by franklyn

More scientific brilliance from the resident retard. And so well formatted too!

woelen - 13-8-2015 at 03:03

Quote: Originally posted by franklyn

To be more specific, electrolysis of a solution of NH4Cl with a platinum anode or graphite anode. At the anode, NCl3 and HCl are formed. At the cathode you get H2 and NH3.

Avoid mixing of anode and cathode liquid. For small (demo) amounts, preventing mixing is not really necessary. You have to stop the process anyway if you want to stay safe.

@franklyn: Why so terse?

[Edited on 13-8-15 by woelen]

The Volatile Chemist - 16-8-2015 at 04:47

I do have two graphite rods, is any high amperages needed? I'll have to work on a way to separate the solutions so it doesn't decay instantly.

franklyn - 16-8-2015 at 09:13

The quiddity (click) of what to search for — Googling NCl3 electrolysis , got me these. My cost 5 minutes.

http://nitrogen.atomistry.com/nitrogen_chloride.html

http://books.google.com/books?id=YtE5AQAAIAAJ&lpg=PA680&...

Revealing the mechanism of indirect ammonia electrooxidation
Attachment: Revealing the mechanism of indirect ammonia electrooxidation.pdf (644kB)
This file has been downloaded 1071 times

Obtained Googling NCl3 oil
http://www.lateralscience.co.uk/oil

http://projectseminars.org/report-risk-and-control-of-nitrog...

Instant gratification
http://www.youtube.com/watch?v=9A9Fg-hJy-4

Old bones , Duh
http://www.sciencemadness.org/talk/viewthread.php?tid=2079

S.C. Wack - 16-8-2015 at 09:15

Good thing mbrown3391 didn't use a flash for his photo. His last login coming 2 weeks after his last posts starting a cyanide thread is duly noted; well done.

Roscoe and Schorlemmer suggests clamping a large FBF full of Cl above a lead saucer in a bowl of a warm saturated solution of NH4Cl, in the hope that NCl3 that collects on the flask all drops down and is collected.

Those who don't want to blow things up prepare solutions and use them right away. Inorg Syn used carbon tet with chloroform. Hentschel via JCS Abstracts used benzene:

3000 c.c. of a solution of bleaching powder, containing 22.5 grams of active chlorine per 1000 c.c., is gently agitated in a stoppered flask of 5000 c.c. capacity, and cautiously treated with a 10 per cent. solution of hydrochloric acid until a portion of the solution no longer yields gas when mixed with excess of a 20 per cent. solution of ammonium chloride. About 300 c.c. of the dilute acid is usually required for this purpose, and when the proper quantity has been added, the liquid is treated with 300 c.c. of a 20 per cent. solution of ammonium chloride, and then vigorously agitated with 300 c.c. of benzene during 30 seconds, the liquid meanwhile being protected from light; the benzene is separated from the aqueous liquid, and transferred to a folded filter containing 20 grams of crushed calcium chloride.

The Volatile Chemist - 21-8-2015 at 12:10

OK, have some procedures to try, as soon as I finish up my chemistry summer work :/

specialactivitieSK - 23-8-2015 at 23:28

https://www.youtube.com/watch?v=kgvEn7l-ZZk

https://www.youtube.com/watch?v=9A9Fg-hJy-4

The Volatile Chemist - 30-8-2015 at 13:42

I'd seen both of these videos. The method in both of them doesn't appeal to me much, as I prefer small aqueous amounts free of as many ligands as possible. I have a few experiments in mind that'd require such conditions.

Velzee - 21-9-2015 at 11:55

Quote: Originally posted by S.C. Wack

Good thing mbrown3391 didn't use a flash for his photo. His last login coming 2 weeks after his last posts starting a cyanide thread is duly noted; well done.

Roscoe and Schorlemmer suggests clamping a large FBF full of Cl above a lead saucer in a bowl of a warm saturated solution of NH4Cl, in the hope that NCl3 that collects on the flask all drops down and is collected.

Those who don't want to blow things up prepare solutions and use them right away. Inorg Syn used carbon tet with chloroform. Hentschel via JCS Abstracts used benzene:

3000 c.c. of a solution of bleaching powder, containing 22.5 grams of active chlorine per 1000 c.c., is gently agitated in a stoppered flask of 5000 c.c. capacity, and cautiously treated with a 10 per cent. solution of hydrochloric acid until a portion of the solution no longer yields gas when mixed with excess of a 20 per cent. solution of ammonium chloride. About 300 c.c. of the dilute acid is usually required for this purpose, and when the proper quantity has been added, the liquid is treated with 300 c.c. of a 20 per cent. solution of ammonium chloride, and then vigorously agitated with 300 c.c. of benzene during 30 seconds, the liquid meanwhile being protected from light; the benzene is separated from the aqueous liquid, and transferred to a folded filter containing 20 grams of crushed calcium chloride.

You think he died(or was injured)?

S.C. Wack - 21-9-2015 at 12:36

Odds are about even that, or he made it to appear so, or his chemical adventures aroused interest from the authorities, by one of the several ways that could happen. His last posts were saying he was going to make HCN...maybe he used a flash for a NCl3 closeup instead. It looks like he's not logging in any time soon, so we'll never know, that's the well done part.

Sakomekieh - 12-12-2015 at 17:01

I tried the synthesis using muriatic acid, sodium hypochlorite and ammonium salt(In this case Ammonium Bicarbonate), but no NCl3 formed... First i reacted ammonium bicarbonate with excess muriatic acid, then mixed the solution with hypochlorite. The solution immediately turned green, but no NCl3 was formed.

I think that organic impurities in muriatic acid decomposed the NCl3(the smell of chlorine was really strong), but i'm not sure :mad:

The Volatile Chemist - 14-12-2015 at 15:09

Quote: Originally posted by Sakomekieh

I tried the synthesis using muriatic acid, sodium hypochlorite and ammonium salt(In this case Ammonium Bicarbonate), but no NCl3 formed... First i reacted ammonium bicarbonate with excess muriatic acid, then mixed the solution with hypochlorite. The solution immediately turned green, but no NCl3 was formed.

I think that organic impurities in muriatic acid decomposed the NCl3(the smell of chlorine was really strong), but i'm not sure :mad:

Eh, I tried it too. It's also possible the concentrations simply weren't strong enough. I had the exact same thing happen to me this summer.

PHILOU Zrealone - 15-12-2015 at 10:53

Better switch to more stable chloramines... see Axt tread about energetic haloamines... alkylic dichloroamines and dibromoamines...
CH3-NCl2
Cl2N-CH2-CH2-NCl2
Br2N-CH2-CH2-NBr2

The Volatile Chemist - 16-12-2015 at 18:03

Why? Certainly if I wanted the NCl₃ for an explosion, I would. But I want it for experimentation.

PHILOU Zrealone - 18-12-2015 at 11:45

Quote: Originally posted by The Volatile Chemist

Why? Certainly if I wanted the NCl₃ for an explosion, I would. But I want it for experimentation.

What experimentation?
If you want to use it for safe experimentation it must be:
1°) In extremely low quantity (<1 drop)
2°) In diluted form
a) into a safe unreactive solvent (CCl4?)
b) into an inert dilutant gas
Manyfold of course!

The Volatile Chemist - 30-12-2015 at 14:40

You know that's not what I'm planning, but I'm also not planning experimenting with more than a drop. Yes, it's explosive, but as long as I do the experimentation outside, with physical protection, it doesn't seem like the other three things in your list are necessary. Just my opinion, for me, though.

PHILOU Zrealone - 31-12-2015 at 09:40

Even a drop in glass, hard plastic or metal, may cause cutting or piercing schrapnels...

NCl3 has the bad habit to detonate upon contact with dirt, grease, almost anything can get a catalytic effect...

So take care of you with your physical protection.

The Volatile Chemist - 31-12-2015 at 11:19

I will, though I doubt I'll get to look at it soon, as I have a lot of other things on my plate. I'll do it in the summer, outside, with protection and such. If I do it at all.

aga - 31-12-2015 at 12:30

Quote: Originally posted by The Volatile Chemist

I'll do it in the summer, outside, with protection and such. If I do it at all.

"Evolution has many mechanisms to eradicate unsuitable genetic traits from a species [1].

This mostly involves the carrier of those genes dying before breeding [2].

It can also be accomplished by disfigurement, including removal of crucial components of the breeding apparatus [3]"

[1] God, 0
[2] Darwin, 1869
[3] aga et al 2015

[Edited on 31-12-2015 by aga]

The Volatile Chemist - 31-12-2015 at 15:09

You speak with experience. I respect that...not really.

'
'We'll see what happens' is a fine trait, thank-you-very-much.

Sakomekieh - 4-1-2016 at 15:11

Quote: Originally posted by The Volatile Chemist

Quote: Originally posted by Sakomekieh

Eh, I tried it too. It's also possible the concentrations simply weren't strong enough. I had the exact same thing happen to me this summer.

I think you are right. The muriatic acid was too weak, it doesn't even attack aluminium vigorously.

Someday i will try using pure conc. Hydrochloric acid(35%), ammonium bicarbonate and bleach.

The Volatile Chemist - 6-1-2016 at 15:45

Well, my muriatic acid is ~35% conc., just has a lot of iron in it.
I'll get to purifying it when I have a torch to bend glass and make a bubbler.