Sciencemadness Discussion Board


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Al Koholic - 23-11-2003 at 14:36

By mixing 500ml of 4M NH3 solution with 250ml 3M NaOCl solution a colorless, warm solution was obtained with the production of heavy, thick, white fumes in the flask. The fumes smelled relatively ammoniacal but also had a distinct sweetish component.

Next, to this warm solution was added 155ml of 30% H2SO4. The addition was accompanied by much of the same white fume production and a significant warming of the flask. During the addition, much effervescence was noted with lots of bubble production. Towards the end of the acid addition, the bubbles kept coming out but the white fumes disappeared along with the smell of ammonia. The now hot solution smelled very strongly bad. Thats about the best I can think of to describe the smell. I think this might be just smelling hot sulfuric acid vaporizing out of solution as the solution is now quite acidic.

Anyway, after boiling down to 1/3 volume and cooling there should be a crop of hydrazine sulfate. The only weird part about this whole procedure was the evolution of the white fumes and the bubbling during acid addition. Shouldn't the acid addition produced warming but no bubbling as all that should happen is neutralization of the ammonia and hydrazine salt formation?

unionised - 23-11-2003 at 14:57

Wot? no glue?
2NH3 +3NaClO ---> 3NaCl +N2 +3H2O
Unless you are rather lucky and add suitable materials to prevent the parasitic reactions.

Al Koholic - 23-11-2003 at 16:15

Yeah, I decided to omit the gelatin component from the reaction at the last minute. I am partly hoping that it will still work because I used all very pure chemicals, and partly just trying it out without gelatin to see what kind of yield results. If this doesn't work, then I suppose I'll have to consider using the gelatin next time.

It's easy enough to prepare the solutions so no big deal if this turns out bad.

unionised - 23-11-2003 at 16:20

Good luck, I seem to recall that a few ppm of metals will kill the reaction and that temperature control is very important.

Qualitative & Quantitative Test for Hydrazine (Sulphate)?

chemoleo - 23-11-2003 at 16:32

Well, I just tried to make some hydrazine sulphate according to my post in the hydrazine oxidiser exploration thread ( ).
It appears that I succeeded making it, the precipitate I get however (after adding LOTS of H2SO4) does not seem that crystalline, it's more fluffy and unstructured?!? (without a microscope tho, I guess it's hard to tell)
I used analytical reagents, including destilled H2O, commercial labgrade NaOCl, NH3, and Na-EDTA as a chelator.

Anyway, now that I have some precipitate, I will have to analyse it to make sure that it is the real stuff.

So, is anyone aware of a *specific* test for hydrazine, a test that won't be fooled by any of the (im)possible reaction products, such as NH4Cl, breakdown products of EDTA (which would be some sort of amines), etc etc?
I am thinking of some complex formation test, i.e. with metals from the transition series.
Even better if this would work in a titratable manner, where one could measure the concentration in solution rather than weight...
Anyone got something on that?

Thanks, Chemoleo.

PS:Unionised - where did you hear that from , about the importance of temperature control?

[Edited on 24-11-2003 by chemoleo]

unionised - 24-11-2003 at 14:26

Silver mirror test and Greenwood and Earnshaw.

Marvin - 24-11-2003 at 20:34

Al Koholic,
I wouldnt bet much on your chances.

The Rashig synth requires a *large* molar excess of ammonia for reasonable yeilds. 10:1 or 15:1 are not unheard of and these yeilds can reach 35% ish.

Temperature control is required according to the synth by Mr Anonymous with urea.

It is not required for the ammonia synthesis. Mixing of the liquids cold then heating as rapidly as possible, followed by boiling off 2/3rds of the solution are normal and give the best yeilds typically. Since this liberates all the excess ammonia, I would imagine this is not a pleasent process.

Oxidation tends to be a bit nonstoichiometric but if you use hyperchlorite, you should rapidly have it oxidised to mostly nitrogen with only a little ammonia. Should give a reasonable idea if you have what you think you have. If you use excess hyperchlorite, oxidation of ammonia to nitrogen shouldnt be a problem.


Al Koholic - 24-11-2003 at 23:48

Well are on the money again...

I just boiled the solution down from 1000ml to 300 ml and put it in the firdge. After cooling to about 2C, nothing had crystallized...

I must have assumed that my "very pure" chemicals were purer than they really were.

In my next trial I will use gelatin to prevent this problem and will heat the solution more rapidly than I have been.

Edit: Just thought I would ask you all but what are considered "typical" yields if the reaction goes well? 50%? Worse? I get the impression that it won't be good...

[Edited on 25-11-2003 by Al Koholic]


KABOOOM(pyrojustforfun) - 26-11-2003 at 20:27

<b>copper dihydrazinium sulfate</b>
&nbsp;Properties: Bluish powder; mp above 300°C, starts
&nbsp;&nbsp;to decompose at 140°C; very slightly solouble in
&nbsp;&nbsp;watrer, 250 ppm @ 80°C.
&nbsp;Hazard: Moderately toxic by ingestion or inhalation.
&nbsp;&nbsp;Skin and eye irritant.
&nbsp;Use: Foliage fungicide.

addition of copper sulfate to hot HS soloution (HS is very solouble in hot water) should produce it. you can then weigh the precipitate ...


fritz - 30-11-2003 at 03:19

For the preparation of Hydrazine-sulphate from Hypochlorite and Ammonia solution there should no excess of chlorine in the hypochlorite solution. So prepare The Hypochlorite-sln. by bubbling Chlorine through an excess of NaOH(100ml 2n solution + 6g Chlorine). I think Ican remember reading somewhere that EDTA isn´t a good idea in the production of Hydrazine. Try Calciumhydroxide (think I got this from FEMFEP) this would do nice.
After evaporation precipiate the Ca as sulphate at pH~7 filtrate an precipiate the Hydrazine-Sulphate

Anhydrous Hydrazine synthesis, Unabridged for Newbies, Kewls and the other downtrodden misunderstoods...

Hermes_Trismegistus - 1-12-2003 at 17:52

I found the following sythesis on Megalomania's site, being a Science Newbie (although eager) it helped me to more fully understand the reactions that were occuring, as well as being written in plain language.

Synthesis: Prepare a solution of 1500 mL of 28-29% ammonium hydroxide, 900 mL of water, 375 mL of 10% gelatin solution, and 1200 mL of normal sodium hypochlorite solution. It is absolutely imperative to use distilled water, the presence of any contaminant ions will screw up this reaction! It is possible to use starch, glue, or glycerol instead of gelatin, but they are inferior. Mix these chemicals in a large glass dish, like a pie plate or bowl, or just use several portions, as this is nearly a gallon of liquid. This mixture is heated as rapidly as possible and boiled down to one-third of its original volume. The solution is then cooled thoroughly with ice and suction filtered twice to remove any impurities. When filtering, first use towels (like a washcloth), then use regular filter paper on top of some cloth (like from a T-shirt).
The resulting liquid is dilute hydrazine hydrate. To make concentrated hydrazine hydrate, mix 144 mL of dilute hydrazine with 230 mL of xylene in a round-bottomed 500-mL Florence flask. Fractionally distill the mixture in an atmosphere of nitrogen, the xylene will first pass over with most of the water, then the hydrazine will pass over. Keep the fractions separate of course. The resulting hydrazine hydrate will be 90-95% hydrazine. This concentration procedure is meant for 60% hydrazine hydrate, since the hydrazine hydrate prepared above may be greater or less than 60%, some experimentation may be needed to find the proper amount of xylene to use (more xylene is needed for dilute hydrazine, less for more concentrated hydrazine).
To obtain anhydrous hydrazine, mix 20 g of potassium hydroxide per 100 g of >90% hydrazine hydrate in a beaker, let this mixture stand overnight so much of the water can be withdrawn. After standing, filter the solution to remove the hydroxide. Add to the filtered liquid an equal amount by weight of sodium hydroxide. Place this mixture in a round-bottomed 500-mL Florence flask, reflux for 2 hours, then distill in a slow stream of nitrogen. You must use nitrogen, distillation in air may lead to an explosion!

Success with hydrazinium sulphate

chemoleo - 1-12-2003 at 18:20

I made some hydrazine sulphate according to my post in the hydrazine oxidiser thread.
Noteworthy observations:
-when mixing the ammonia with the NaOCl (both at 4 deg C), no bubbles form. However, slight heating to 40 deg or so produces lots of bubbles, that subside quickly.
-boiling it initially keeps the temp at around 75 deg, until after 1/4 or so are evaporated, then the temp rises to 95 deg.
-at 1/4 of the original volume, crystals appear, which were filtered off. I hypothesise this is NaCl or NH4Cl.
-Using EDTA, the solution is slightly yellowishh at that point, with 2 g pig skin gelatine the solution is thick and a nice chlorine grean (?!?)
-Addition of H2SO4 to the solution at 4 deg cause massive gas evolution, and at least in part contains chlorine (smell)
- Temperature heats up considerably, to 50 deg. This was cooled again, and more H2SO4 added. Preciptiation (fairly fluffy) finally ensued. pH at that poiint is well below 0. Gas continued to evolve, which was green (not much annymore thoug)
-This was spun down (centrifuge) and the chlorine-smelling pellet was redissolved in hot water, and recrystallised.
-Due to the extreemly low pH at even that point, the preciptant was suspended with 96% ethanol, filtered, and repeated. Final product is a white crystalline powder that is not much soluble in water.
-Gelatine did indeed a higher yield than EDTA, even though there wasn't much to start off with. Hell, got a few grams, not like kilgorams I dreamed of :(

Addition of that to CuSO4 did indeed yield a turquoise blue precipitate.

What I am wondering...Why does gas evolve once the solution is boiled (i.e. at 40 deg) - gelatine or EDTA should prevent the latter!
I used analytical reagents of course. Anyway, this definitlely decreases my yield!
Why is there gas evolution upon the addition of H2SO4? Especially chlorine??? This one eludes me. NH3 is in excess (big time), so all of the NaOCl should react....

A definite peculiarity...

Al Koholic - 2-12-2003 at 20:16

As a matter of fact, I ran into exactly the same problems. The bubbling upon supposed neutralization of the hopefully hydrazine solution is most perplexing. I did not notice a green gas but more of a thick cloudy white vapor of some kind during my neutralization. I also used a somewhat large excess of NH3...

Lets think about this...I did not use gelatin or EDTA and still I get the bubbling upon warming of the solution and neutralization. You used both gelatin and EDTA and notice similar bubbling although noting the presence of chlorine by smell which I did not notice. Perhaps some side reactions are occuring in both of our cases with different outcomes? Seems unlikely...

Did you notice a really bad smell when you added the H2SO4 to the solution for neutralization? It was bad in the sense that rotting food is bad...nauseating...not like the pungent sting of Cl2.

[Edited on 3-12-2003 by Al Koholic]

chemoleo - 2-12-2003 at 20:34

hmm, the weird thing was that H2SO4 addition generated green bubblkes.... hence thinking its chlorine. No idea why the NaOCl hastn reacted completely at that point. This wasnt much, compared to the LARGE amount of bubbling caused (with no colour) as soon as I started adding H2SO4 (the white fumes/vapour you mentioned - definitlely not caused by acid/water mixing) - the green bubbles (not very much admittedly) only started appearing after I had already added quite a bit of H2SO4!
I don't think it's down to the complexing agent (gelatin or EDTA) as I got the same result for each. It must be a generic problem.
Yes, and when I added H2SO4, besides all the gases it produced, the smell it produced didnt seem healthy, thats why I kept it under the fume cupboard at all times. the smell was rather asphyxicating that is. Kind of not painful, but pleasant either.
Anyway, I got about 4 grams hydrazine sulpate now. I am sure it's the right stuff, but still a terrible yield. Its a terrible yield considering I started off with several molar NH3, and nearly 400 ml total (sorry I havent calculated the theoreticl yield yet)
I wonder how they do this industrially, with yieldslike that it couldnt have been ever economical!

[Edited on 3-12-2003 by chemoleo]

Al Koholic - 3-12-2003 at 12:30

Chemoleo...what concentration sulfuric acid did you use to neutralize the mixture? I used 30-35% battery electrolyte.

Also...I did notice just now that we obtained very similar gas production even though I neutralized my mixture before reducing the volume by boiling.

I am thinking about some experimenting in the near future where I will make more NaOCl, NH3, and use gelatin this time. I will boil this mixture down to 1/3 the original volume or so and then divide it up into different portions. I will test different concentration H2SO4, different concentration HCl, AcOH, etc. This should help us determine if the phenomenon is specific to sulfuric, or to acid neutralization in general. Of course I'll be noting smells, appearances, quantities, etc... Other than that I am at a loss for why that gas production occurs and I can't shake the feeling that the yield is being affected by the neutralization.

I also found today that my test definetly did not produce any hydrazine. I added some copper sulfate solution to my 250 ml of boiled down solution and did not get a turquoise precipitate. It is good to know about that test however.

More to the production of Hydrazine

chemoleo - 5-12-2003 at 17:11

ok here's the answer:
I used 98% H2SO4, and I would do the same if I were you. Else you will just dilute the putative hydrazine solution needlessly. Boil it down to less than 1/3 of the original and salt crystals will appear. Filter them off, thats already a preliminary purification step. Then add H2SO4, and lots of gas will be produced, plus temp goes up. Cool again. Then add more H2SO4, and a fluffy product appears. Dont add too much because some of it seems to redissolve the more H2SO4 you add. maybe because the bisulphate is produced? I know thats the second form of hydrazinsulphate, but I don't know about it's solubility.
Anyway, collect the precipate and recrystallise once or twice in hot H2O.
After this, the solution is still strongly acidic, which is in part due to the H.S., but also due to left over H2SO4.
For this I mixed the recrystallised H.S. with a large excess of 96% ethanol, it resuspends nicely. If you do that twice, you can dry the final pellet on air, and it is NOT particulary hygroscopic, unlike the crude product that wasnt ethanol-washed!
I do think the gas evolution is specific for strong acids, such as H2SO4. It probably won't havppen with HAc.
Strange I also cant shake the feeling that neutralisation somehow affects yields...trying to think of ways to avoid this, to isolate hydrazine without using steps such as destillation!
Ideas anyone?

PS By the way, I also had the very strong impression that boiling the neutralised/strongly acid solution to decrease volume & increase yield doesnt help, doing this seems to destroy the H.S. as yields are pathetic after a second boil!

[Edited on 6-12-2003 by chemoleo]

BromicAcid - 5-12-2003 at 19:30

First and foremost I found something odd in one recently aquired library book, it mentioned that the resulting mixture from the production of hydrazine can be distilled instead of being precipitated, it said, and I quote "Distilling the resulting dilute solution of hydrazine yeilds a 58.5% hydrazine - 41.5% water azetropic solution..." (Inorganic Chemistry of Nitrogen, Jolly, 1964) which of course can be concentrated by distillation over sodium hydroxide pellets. Sounded interesting to me but I would have to look up the azetrope distillation temp.

But secondly, there is a book, it's new, and someone out there in college might have it in their library: "Hydrazine and its Derivatives : Preparation, Properties, Applications" this book came out in 2001 for the second edition and it is 2232 pages! :o Anyways, I had the first editon at another local library and ordered it, but here is what caught my eye, over one hundred pages on the preparation of hydrazine using numerous methods (Im guessing some interesting curiosity methods) regardless, amazon has it, but it's $393.12, I mean, come on! I'm not made of money!

chemoleo - 5-12-2003 at 19:59

by the way, I think I figured out why this green chlorine- like gas evolves once the solution is *strongly* acidified: NaOCl disproportionates to NaClO3 (!!) and NaCl. As I didnt use freshly prepared NaOCl, I am sure the green gas was ClO2, which is liberated when HClO3 is formed....
Lovely :)
Thank the gods of chemistry for a superb fume cupboard!

Al Koholic - 6-12-2003 at 08:24

Which is exactly why I didn't get any green gas because my solution had been freshly prepared!!!

BromicAcid - 12-12-2003 at 19:45

Okay, I checked out "Hydrazine and its Derivatives" a few days ago and boy oh boy is it a good read. For preparation of hydrazine the highlight would have to be the Bergbau-Bayer-Whiffen Process involving the formation of a ketazine as an intermediate step to boost hydrazine yeilds, over 100 pages on how to make it so I will look over thoughly and condense out all the pertininent information.
But the most interesting thing I found so far was

"If the carbonyl group in urea is removed, the remaining two amino groups are likely to join and form hydrazine. Accorind to a patent by Passino [Passino, H.J.: Manufacture of Hydrazine, U.S. Pat. 2717201 (6 Sep 1955), M.W. Kellogg Co.; CA 50, 2131b.] this can be acheived by heating urea with nickel or another carbonyl-forming metal. Under the conditions of the reaction, the metal carbonyl decomposes such that the metal acts as a true catalyst, and hydrazine and carbon monoxide are the only products:

H2N-CO-NH2 --Ni--> H2N-NH2 + CO

Sounds pretty interesting, I will post one more reply here when I finish this book, it's a great read!

unionised - 13-12-2003 at 03:26

It certainly sounds interesting but I have a nasty feeling it won't work. Partly because I don't think the 2 C N bonds would break like that, but mainly because Ni would catalyse the decomposition of the hydrazine. Ni(CO)4 is stable up to about 250C so it would need to be hotter than that for the Ni to be regenerated. I'm not sure that N2H4 would survive in the presence of a good dehyrogenation catalyst at 250C.
This may have been patented, was it ever used? If it were, it would be a vastly cheaper way of making hydrazine than NaClO oxidation.


KABOOOM(pyrojustforfun) - 13-12-2003 at 20:59

pleaaaaase don't tell me you don't have scaner. let's make us indebted to you for a whole life!

btw I read in chemical dictionary that Ni(CO)<sub>4</sub> explodes @ 60°C (I think it means it decomposes explosively. not much different though!) it boils @ 45°C it's a known carcinogen, tolerance: 0.05 ppm in air

BromicAcid - 14-12-2003 at 09:07

Okay, so I've never looked for a patent before. I didn't know you could just go to: and search the patent number and get pictures of the pages of the patent.

The reaction of urea with nickel is quite interesting, run between 40 C and 130 C with nickel powder present at between 2% and 50% by weight it seems to work better as the percent of nickel gets closer to 50%. It says that iron can be substitued at higher temperatures. The patent doesn't mention yields but it tells what reaction products to expect at what temperatures.

If I had to guess why this method was not used industrially I would have to say high cost of nickel as a material that might be consumed and possibly the reaction might be explosive if scaled up too much. It just seems to good to be true, and simple.

unionised - 14-12-2003 at 12:49

I can only think of one reason for not citing the yield of the reaction, and it's the same reason that nobody uses this method.
BTW, Ni(CO)4 is stable at 40C unless there is air present.

Marvin - 14-12-2003 at 16:12

Ive known about this 'wonder' method of hydrazine synthesis for some time. Actually about 8 or 9 years and I have paper copies of both patents. On the face of it, it looks looks good. Anhydrous hydrazine in 1 step from OTC chemicals.

Aside from the possible worry of the TM metal causing the hydrazine mixture to explode there are more deeply worrying problems. Nickel carbonyl is ungodly toxic. The general industry allowable exposure for nickel carbonyl is 10000 (ten thousand) times lower than for hydrogen cyanide (OSHA PEL). That speaks volumes to me.

Iron carbonyl is probably a little less toxic and iron powder is more easily available but one of the series of iron carbonyls boils at a very similar temperature to the hydrazine produced and I wasnt able to find out more more about it. I gave up trying to find a way to make this reaction safe enough in my head to try out for real but there are plenty of astrolite worshipping morons out there that would probably not come to the same conclusion. Hense keeping the information quiet.

BromicAcid - 6-3-2004 at 11:54

I've collected everything that I need to perform the reaction of urea with nickel powder to produce hydrazine. All new glassware (50 ml 24/40 RB sure looks sized odd), good big stir bar, nice new hotplate, liebig condenser, etc... I'm going to wrap the whole apparatus in two or three layers of aluminum screening in case of detonation for some reason or another.

However I'm wondering about the exit gasses, CO probably mixed with some Ni(CO)4, my current plan is to run them into a Bunsen burner and subject them to high temp, burn the CO and the Ni(CO)4. But does anyone know of a solution I could run them into to neutralize them, seeing as how that would be an easier alternative. I know that a NaOH solution will react with CO but I believe it's only at elevated pressure.

One other thing, Marvin you said you have paper copies of both patents, I only know of the one patent, what's the patent number for the other?

And one last thing, I was considering running the reaction in a solvent. Just preliminary solubility queries yield urea to be soluble in benzene and hydrazine to be pretty much insoluble. So with efficient stirring the hydrazine should sink to the bottom and the nickel will be in better contact with the urea in solution, also the carbonyl will be soluble in the benzene so it might help to moderate the reaction. But I've got a bad habit of trying to modify reactions that I've never run so I'm going to run it by the book a time or three but the solvent sounds like a reasonable modification to me.

[Edited on 3/7/2004 by BromicAcid]

unionised - 7-3-2004 at 09:18

Urea is nearly insoluble in ether and chloroform. My guess is that it wouldn't disolve well in benzene. I can't find a lot of solubillity data on hydrazine but it is soluble in the lower alcohols and in water. I would be a little suprised if it isn't reasonably soluble in benzene. (Ammonia is, ethylene diamine is too).
I would also be wary of using benzene as a solvent if I could avoid it, on the grounds of toxicity. Then again, I don't supose you will be sniffing the reaction mixture.
IIRC you can use solutions of copper (I) compounds in ammonia to trap CO. NaOH won't do the job under normal conditions (About a million years ago I did a university practical investigating the CO/CO2 +C equilibrium. The CO was measured over NaOH soln)

BromicAcid - 7-3-2004 at 09:32

I don't know how soluble it is but under the entry for urea it is listed as being soluble in benzene. The main reason I was considering a solvent is that looking though the patent it states that at low temperature reaction 40 - 60 C a large quantity of nickel carbonyl vapor comes over. I figure with an organic solvent not only will the nickel carbonyl be absorbed and hopefully have more time to decompose before leaving the solution, but it will afford better contact between the nickel catalyst and the urea.

So I'm looking for a solvent with the following properties:

1) Urea is atleast slightly soluble in it
2) Nickel carbonyl is soluble (it's soluble in most organics anyway)
3) Bp >60C but, <100C unless, >135C
4) Does not form azetrope with NH2NH2
5) Will not react with hydrazine, finely divided nickel, nickel carbonyl, urea, or carbon monoxide

By the way, does anyone have solubilites for nickel carbonyl in any solvent besides water?

[Edited on 3/7/2004 by BromicAcid]

Marvin - 7-3-2004 at 12:55

You do not want to increase the residance time of the nickel carbonyl in the reaction mixture. It wont decompose at the low temperatures so the reaction proceeds forwards mainly by loss of nickel carbonyl (the reason nickel needs to be in excess)

This reaction is reversable, that is to say CO and hydrazine will combine in the presence of nickel metal to form urea (mislaid my reference for this, but given the nickel process for hydrazine production, I didnt need much convicing that it was reversable). More of a pain is that hydrazine left in the reaction mixture will react with urea to form semicarbazide, which constitutes a loss if you only want hydrazine. You can form urea/semicarbazide in the reciver if you use iron salts/other metals as explained in the patent.

Hydrazine hydrate attacks glass, but Ive been told dry hydrazine does not. So long as the mixture is fairly water free, this might not be a problem for you. A more worrying rumor is that ground glass joints can cause hydrazine vapour to ignite/detonate. I dont remeber where I read this, or if this is accurate.

If you are going to try the process, I would avoid the low temperature method producing large amounts of the nickel carbonyl and use the temperature neer the decomposition point of urea yeilding almost entirly carbon monoxide, and condensing the hydrazine out of this.

Call me daft, but I think it might be preferable to mix the CO with the fuel of a gas burner, flush the system continually with say butane and burn the offgas. Pass the gas through the cold setup in order to exhaust the oxygen for some time before igniting it (you know this). Butane should be resonably inert to the reaction.

Other patent number is 2675301. It seems almost identical. Its actually mentioned in the later paper.

unionised, remeber that ten thousand factor I mentioned earlier? I really dont think benzene would make the slightest difference to the overall risk. I do think it would be detrimental to the experiment though.

BromicAcid - 7-3-2004 at 13:09

Marvin, I was unaware that hydrazine and carbon monoxide in the presence of nickel could form urea, thanks a lot, you saved me on that one. I read that other patent, I missed it in the first place because I copied the other patent onto a wordpad document by hand and just ignored the header.

Supposedly according to it the reaction of urea with hydrazine is minimized at atmospheric pressure and nickel carbonyl losses are significantly less at elevated temperatures. Oh well, I guess I'll run this one in an oil bath.

I was somewhat worried about the explosive decomposition of hydrazine on ground glass joints, supposedly the greatly increased surface area can cause it to catalytically decompose. I'm going to go heavy on the silicon grease and hope for the best, it seems to be a rare phenomenon anyway.

Also I was reading up on hydrazine as a mono-propellent in rockets. It is decomposed by and Ir-Pt catalyst, or a silver screen catalyst, or a tungsten oxide catalyst, or an Ir-Al2O3 catalyst. No mention of a nickel catalyst which was one thing that I was worried about, expecially after seeing the kind of heat and energy generated.

So I guess a solvent is not the way to go on this one, heat and hope for the best, ultra-clean glassware, greased joints, exit gasses run into the intake on a bunsen burner. Wrap in screen as a last ditch save attempt. Magnetic stirring, I think I've got this trial run all figured out. I think I'll go for a charge of 25 g of urea and 12.5 g nickel, bearing in mind that 50% nickel by weight gave the best yields in the patent.

[Edited on 3/8/2004 by BromicAcid]

Marvin - 8-3-2004 at 21:30

Hydrazine is usually distilled on a lab scale in silver retorts. In so much as I had any 'good' ideas when reading this stuff, I wondered about protecting glassware with the silver mirror reaction. It would probably flake off neer joints, but it could help the glassware from cosmetic damage, and also a 'frosted' effect that might catalyse decomposition.

Probably not something youd want to try on a first run, but I think I'm better mentioning it than not.

I see you seem to be avoiding my idea of running the fuel through the equipment and out into the burner in favour of feeding the CO into the air intake of the bunson. As long as you purge the container first I dont see a problem. Other than the toxic compounds, the possability of hydrazine explosions and flying glass that is. Mmm. Did I mention how glad I am that I'm over here and your all the way over there. I'm sure I did.

BromicAcid - 9-3-2004 at 07:35


I see you seem to be avoiding my idea of running the fuel through the equipment and out into the burner in favour of feeding the CO into the air intake of the bunson. As long as you purge the container first I dont see a problem.

Well I didn't want to run fuel through it for two minor worries. One, the propane or whatever lower hydrocarbon I put through the system may well be soluble in hydrazine and contaminate the product to a high degree. And two, running the fuel through the system may increase the pressure because it will be running through the small bunsen burner hole and an increase in pressure in this reaction is associated with a loss of yields.

Both are minor reasons but still things that I've been thinking about. Also, can't hydrazine decompose catalytically without oxygen if the temperature is fairly high. E.g. a flame could flash through the equiptment from where the gasses are being burned off and detonate everything regardless of it being purged of air? If so, maybe I should run the exit gasses through some bath first to prevent it from flashing back into the system.

Did I mention how glad I am that I'm over here and your all the way over there. I'm sure I did.

Not directly but it has been implied. Based on your warnings I have decided to scale down my reaction vessel. Going to run it in 14/20 glassware if I can get one more piece. Wash all the glass with alcoholic NaOH beforehand, oven dry, assemble quickly. Wear proper protection, long gloves, face shield. I have all the equiptment I need to do this safely. But I still won't be doing this for a month or so.

[Edited on 3/9/2004 by BromicAcid]

Smaller scale first?

Polverone - 9-3-2004 at 12:38

I wonder if the form of the nickel powder is especially important to its reactivity. Will nickel powder formed from zinc dust and aqueous NiCl2 solution work, or do you need to prepare nickel formate and decompose it with careful heating while protected from air?

Before you even attempt isolating the hydrazine, why not try this on a test tube scale? Urea decomposes fairly readily with heat anyway, so bubbles aren't the sure mark of success. It would be nice to first verify that you're producing some hydrazine before you scale up. Will hydrazine decolorize iodine starch test paper? I don't have a qualitative test handy, but I'm sure you can look one up.

Also, second idea: assuming that it works on a test-tube scale, how about passing your exit gases through aqueous H2SO4 when you first scale up? Hydrazine sulfate's useful, and it's a conveniently weighable form (for judging yields), and it's far less hazardous. "Walk before you run" as the saying goes.

BromicAcid - 3-4-2004 at 13:33

I'm going to be running a test tube scale first. I got some flint glass and made it into the shape I needed so that it would exit the 'hydrazine production tube' and go a distance then down into some cold 6M H2SO4.

I've got everything I need but some nickel powder. According to the patents nickel of any shape or form can be used. From the powder used in the small scale to the sizable balls of nickel that would be used in the continuous process mentioned where molten urea is constantly circulated through a bed of Ni. I made some NiCl2 but I need to reduce it somehow. It's still slightly acidic and in solution and I'm trying to reduce it without evaporating all the choking fumes then reconstituting then reducing. I tried Al with no success, zinc made a white precipitate, odd. Sn in HCl made small amounts of Ni powder but nothing sizable, let it go for two hours, maybe next time longer. Any easy reducing agents anyone can think of?

Regardless next weekend is the next time I will be experimenting so that is when I will attempt this process, if I can't get my Ni powder by then I will just go with nickel chips.

[Edited on 4/3/2004 by BromicAcid]

Polverone - 3-4-2004 at 14:17

Zinc should have worked. I had no trouble producing nickel dust from zinc dust swirled in acidic aqueous NiCl2. However, my crude "nickel powder" had a considerable amount of zinc left in it. This I was able to remove by treatment with dilute acetic acid (which dissolved the zinc much more readily than the nickel), leaving me with a fine black dust.

I made an attempt today.

BromicAcid - 10-4-2004 at 12:59

  1. This is the apparatus prior to any heating or anything. The test tube has been charged with about 3 g of urea and 2 g of nickel in the form of filings and small pieces also a stir bar was thrown in. This was in an oil bath that also contained a stir bar to allow for evener heating. The thermometer is also placed in there. The exit gasses passed through a piece of flint glass and into a second test tube in an ice bath filled with 6M H2SO4. Immediately after this picture was taken heating was begun.
  2. Close up on the contents of the test tube prior to heating.
  3. The mixture melted at about 137 C and heating was slowed and it stayed relatively constant at this temp. I held it here for about 20 minutes with very very infrequent bubbles. No precipitate came in the sulfuric acid so heating was increased after this time period.
  4. Over the course of one hour heating gradually increased. This picture was taken when the mixture had reached 200C. It was constantly bubbling and expanding from the frothing. At the lower part of the test tube there was condensation that hardened shortly above the oil line. Then there was a space with no condensation and right at the top it was foggy due to additional condensation. Probably just urea. Meanwhile in the other test tube there was no precipitate. However a crust would occasionally form right were the gasses exited into solution and would break off in a semi-circle and float to the top so the top looked like it was covered in a thin skin, hydrazine sulfate?
  5. It was shortly after the last picture that heating was discontinued entirely. Notice the foggy condensation on the generation test tube. The second test tube is somewhat milky white. This occurred because it got so cold after the H2SO4 addition. It cleared up after I let it sit awhile. Although there was a crust on the top the picture that I took did not turn out.
  6. For the test tube that collected the gasses I added copper sulfate as mentioned further up thread as a test for hydrazine. I got a weird thick skin on the top of the liquid after stirring that looked remarkably like ice. Other then that I had no additional tests. I took the other test tube that had generated the gasses and added water to put some of it into solution. It made a milky white precipitate filled liquid. Copper sulfate produced no effect.


To me it looks like it didn't work too well if at all. Gas evolution was only significant at about 160C below that bubbles were few and far between. The nickel came out blackened. A good sign that something was going on with it but other then that nothing else. I plan on running this again next weekend with iron in place of nickel. And if that even doesn't work one last time with 14/20 lab ware with about 20g of urea, seeing how sluggish the reaction was with nickel shavings and such I believe the next attempt will be completely nickel powder. But as it looks now I don't think barely any hydrazine was formed at all.

[Edit] One other thing, the molten urea turned green towards the end, is that normal? I thought that maybe some nickel carbonyl may have formed and by all accounts I would consider molten urea an organic solvent so it could have went into solution and caused the color change.

[Edited on 4/11/2004 by BromicAcid]

unionised - 12-4-2004 at 13:50

I have a feeling that molten urea would disolve nickel oxides and give a green solution. It might well disolve nickel in the presence of air.
The decomposition of urea without a catalyst is a complex affair. How do you propose to prove the presence of hydrazine rather than biuret, cyanic acid, melamine or whatever?
BTW, is anyone in a position to check out this reference?

[Edited on 12-4-2004 by unionised]

The_Davster - 12-4-2004 at 14:28

Unionised: You cannot see that page? All I can see is the abstract, is this the same for you or can you not see the entire document. Anyway here is the abstractif you cannot see it.
<Quote>: Abstract: The catalytic decomposition of hydrazine is studied with regard to the needs of the airspace industry. Data are presentsed on the rate, stoichiometry, and activation energy of hydrazine decomposition on metals of platinum and iron groups and also tungsten and molybdenum, supported by alumina. the effect of ammonia and water concentrations on the rate of hydrazine decomposition is studied. the difference is demonstrated in the rates of liquid-phase and gas-phase decomposition of hydrazine. Phenomena are discussedrelated to the transition from heterogeneous decomposition of hydrazine to the homogenous process.</quote>

[Edited on 12-4-2004 by rogue chemist]

BromicAcid - 12-4-2004 at 15:01


The decomposition of urea without a catalyst is a complex affair. How do you propose to prove the presence of hydrazine rather than biuret, cyanic acid, melamine or whatever?

Being that the solution was hot enough to volitize any hydrazine formed I lead it through a 6M H2SO4 solution. If any hydrazine was formed in any sizeable quantity I was relying on it reacting with the sulfuric acid and forming the firmilar hydrazine sulfate precipitate. Then I could have done some tests on the precipitate, e.g. saturated a hot solution with it then added CuSO4 like mentioned up thread to see if I got a precipitate from that.

Also, if I didn't go for that and just tried to collect a distilate I could have tried to dissolve sodium metal in the distillate and if it went in with no reaction other then color change then I could have reasonably assumed (considering the distillate would be at stp) that it was hydrazine.

Looking over the evidence again I really don't think the reaction would go well. The carbon really doesn't resemble the carbon in carbon monoxide at all so I would assume it would not as easily form carbonyl compounds, it is lacking the electrons to easily form them. However at the higher temperatures with decomposition occuring anyways the carbonyl may well form. Next run, greater quantities, straight nickel powder, longer reaction run.

Maybe I could run it in aqueous solution in the presence of an acid like HCl at about 80C. This would protonate the nitrogen and give it a positve charge which would draw some of the electron density from the carbon and possibly make it more reactive. Just a hope but oh well, time for some experimentation.

[Edited on 4/12/2004 by BromicAcid]

unionised - 13-4-2004 at 01:44

I could see the abstract; I wanted to know what the full report said. In particular, I wanted to know if nickel was among the platinum group metals that decompose hydrazine. If it is then this method is dead.

In the presence of enough HCl to protonate the urea, the metal catalyst will disolve and the urea will hydrolyse.

From 3 grams of urea, if you got 100% yield you would get about 1.5 g of hydrazine which would give you about 6.5 g of the sulphate.
That would dissolve in about 200 ml of water. If you have about 20ml of dilute acid in which to trap the product then you need at least a 10% yield to get any crystals of hydrazine sulphate.
The formation of crystals is not a sensitive, or specific test.

[Edited on 13-4-2004 by unionised]

hydrazine reaction alternative chelator

kazaa81 - 1-5-2004 at 12:44

What chelators can be used instead of EDTA for hydrazine synthesis by NH3 and NaClO?

BromicAcid - 1-5-2004 at 18:59

Cysteine (which is found in egg whites) could probably be used. I don't have the references handy as to which metal ions are the most detrimental to hydrazine formation (I think they all are to some extent) but the necessity of a chelator can be lessened by using something to 'gum up' the mixture as most synthesis on the web say add some gelatin. However your yeilds from the partial oxidation of ammonia are going to be low by definition due to the side reactions.

[Edited on 5/2/2004 by BromicAcid]

The_Davster - 1-5-2004 at 19:01

Does this mean one can use the "slime" from those kiddie chem sets? :P

Also BromicAcid here is a hydrazine test kit.

[Edited on 2-5-2004 by rogue chemist]

Hydrazine Regurgitation

S.C. Wack - 13-5-2004 at 22:23

heskogen was asking some things in another thread, answers that were in other threads, like this one. This got me looking into my notes for additions here. This contains marginally useful factoids but no syntheses, so turn away or don't. This is mostly well-known (I thought) stuff, intended for the reading-impaired, in case it can ever be found useful. So some of it has been mentioned already. Some of it is just stuff I wrote down and don't know why. Feel free to say its not all practical.
First of all, CV 1, 309. Go there. There are references to patents, etc. in addition to the bad-yielding hydrazine prep. Next, The original patents by Rachig: DE192783, DE198307, US910858. His article: Ber., 40, 4580 (1907). Libraries! Microfilm! Also, GB199750.

Explanation for heskogen/the unaware: The reaction NaOCl + NH3 = NH2Cl + NaOH is very easy even cold. The reaction NH2Cl + NH3 + NaOH = N2H4 + NaCl + H2O is not so easy as it requires heating to boiling or close to it. Much easier is the reaction 2NH2Cl + N2H4 = 2NH4Cl + N2. There is another bad reaction: 3NH2Cl + 3 NaOH = 3NaCl + N2 + NH3 + 3H2O. So its kind of important to get the chloramine reacting with NH3 instead of something else.

Some experimental with molar excess NH3/yield based on NaOCl:
6X/25%, 10X/38%, 20X/50%, 37X/63%, 77X/75%.

Rachig found that an additive was needed. Large amounts of glycerol, sugars, and starch increased yield to 40%. Albumin, casein, and glue gave 60% Glue and a large excess of NH3 finally gave 75%. Some experimenting with Knox gelatin (typical American brand), all using 200 ml. 28% NH4OH and 100 ml. 7.5% NaOCl, in wt. of gelatin in mg.'s/yield N2H4:
0/1%, 3/20%, 10/44%, 50/63%, 100/71%, 500/77%, 1000/88%, 2000/92%, 3000/91%, 4000/91.5%

Some quantitative tests for hydrazine, months late:
N2H4 + HCl + KIO3 = KCl + ICl + N2 + 3H2O. The iodate is added until the iodine color goes away. The HCl conc. is such that it is 3-5N at the end point.
Alkaline KMnO4 gives MnO2, alkaline ferricyanide gives ferrocyanide.
A neutral solution that can contain a very small quantity of N2H4 gives a precipitate of the azine when shaken with benzaldehyde (straight out of the imitation almond bottle?), propionaldehyde, or salicylaldehyde. The precipitate dissolves, however, in excess aldehyde. This is not necessarily bad, if you have a lot of aldehyde that can act as reagent and water insoluble solvent. These azines are even less soluble in cold H2O than benzaldehyde is in warm. There is a synthesis of methylhydrazine via this route in Org Syn CV 2, 395. The aldehyde can be regenerated by acid hydrolysis.
Some qualitative tests: Fehlings gives Cu. Acid chromate gives Cr+3. Of course peroxide, persulfate, chlorate gives HN3.
Allegedly an interesting way to plate glass with Cu: 2g. cupric acetate and 100 ml H20 are mixed with enough NH4OH to dissolve the precipitate. 15 ml. 40% hydrazine hydrate is added and all is dripped onto a glass object, at 60C, until a shiny layer of Cu plates the glass. The glass is washed with hot water and put in a bucket of 60C water that is allowed to cool at its own pace.
Similar things can be done with silver (GB524753) and nickel (DE717547) also Pt and Au. This Ni plating may or may not be helpful with procedures that call for 10% KOH in alcohol and other caustic reactions. Even plastic can be plated from hydrazine and ammoniacal Ag nitrate. I would think that this may be useful for preparing supported catalysts. US1164141 uses this to make a non-Raney Ni catalyst.

Saw acetone mentioned. I read somewhere that the ketazine/acetone/water distillate can be acidified with H2SO4, giving the sulfate and acetone, ref: JACS 51, 3394 (1929). Of course the ketazine can be extracted with organic solvent instead.
Saw urea mentioned. I also have here that the max yield with this in any lab is 60%.

Hydrazine sulfate solubility in water, in deg C/g. per 100 ml H20:
20/2.8, 40/4, 60/8.3, 80/12.6
And at 25C in H2SO4, in g.H2SO4 per L/g.N2H4H2SO4 in 100g. solution:
.5/3, 5/2.7, 26.6/1.5, 49/1, 116/.5

Notes on drying: hydrazine hydrate can be refluxed with an equal wt. of NaOH for several hrs, then distilled at normal pressure to give anhydrous hydrazine. If the reflux is skipped the product is 90-95% hydrazine. If not skipped, a temp of 150 is necessary even though the bp is 113.5C.
85% hydrazine hydrate can be dehydrated by heating with an amount of NaOH equivalent to the H2O it has. KOH cannot substitute. 2 layers separate on heating and the layers contain in % by wt.:
C---N2H4--NaOH--H20 -upper layer:
lower layer:
-from JACS 71, 1644 (1949)

The original method for anhydrous hydrazine: NaOCH3 in MeOH + N2H4HCl = NaCl + CH3OH + N2H4. The alcohol is distilled off, and you must really want anhydrous hydrazine if you abuse your Na in this fashion.
Dilute hydrazine solution has been concentrated by hot (over 50C) anhydrous NaSO4, with the decahydrate crystallizing on cooling. Some water and some sulfate will remain though.
Xylene has been mentioned for azeotropic drying of dilute solutions, and Vogel recommends this to obtain 90-95% hydrazine hydrate. Toluene should work, though not as well.

Last factoid: H2SO4 precipitates, at best, only 90% of the hydrazine in dilute solution.

I'm sure that none of this is true and all the important parts have been left out. You know how books are!

BromicAcid - 27-9-2004 at 11:18

Today is the day, I've completed a fair bit already and here is the step by step:

  1. Placed 8 g of nickel oxylate dihydrate into a 50 ml glass volumetric and stoppered loosely with a 10/30 stopper. Also a very small amount of HgO (<.0005 g) was added, this has a catalytic effect of promoting nickel carbonyl formation.
  2. Heated slowly on a hot plate, first lots of water came off, I took the stopper out to just let the water evaprate away then it started to turn black at the bottom so I re-stoppered it. It continued to release gas until the whole mass turned black. Upon swirling the hot flask the power was free flowing. I set aside and allowed to cool.
  3. I esitimated that my reaction should have yeilded slightly over 2 g of nickel power so while it cooled a little bit I weighed out 6 g of urea and placed that along with a stir bar in a 100 ml 14/20 RB flask.
  4. All parts were previously washed and washed again with distilled water. They were dried in an oven and the joints were wrapped with teflon tape.
  5. I added the cooled nickel powder to the urea flask, it formed little needles on the stir bar. I quickly put the still head in place and the rest of the assembly, a gas entry tube was put into the thermometer hole on top.
  6. H2S (which supposedly increases the activity of nickel toward forming the carbonyl) was slowly put into the system, a very small amount total.
  7. The gas entry tube was replaced with a stopper and the apparatus was lowered into an oil bath and heating begun.
  8. The mixture liquified and the urea took on a blue tint, probably some nickel oxylate that didn't decompose.
  9. The mixture is currently heating away as it has for almost an hour. It is being held at 150 - 160 C and urea is crystalizing on the inside of some the glassware near where the reaction flask is.

There is no distillate so far, but I will keep posted. If I get distillate I will attempt to dissolve sodium metal in it. If it is anything but liquid ammonia (impossible) or hydrazine it should react appreciably.

Edit: Ran the reaction for 1 hour 30 minutes, no distillate, tiny little hairs of urea condensed on the inside of the reaction flask though. Urea crystalized though 80% of the glassware, no liquid anywhere. Heavy ammonia smell even though exit gasses were being washed.

Then I turned up the heat to 200C for 15 minutes, urea decomposed more but no sign of a distillate. :(

I think I'm done trying this for a while until someone else shows some promising results. If freshly made nickel powder will not react with urea with every accomidation made for the formation of carbonyl, I give up.

[Edited on 9/27/2004 by BromicAcid]

Theoretic - 27-9-2004 at 13:18

Hydrazine reacts with Na to form NaN2H3

BromicAcid - 27-9-2004 at 14:24

I've read in a few places that the solvated electron trick can be done with hydrazine as well as liquid ammonia to yeild the characterisic blue solution.

However I'm sure the reaction takes place, similar to the anologous reaction between ammonia and sodium to produce sodium amide, which is catalyzed by trace metal ions and moisture.

Edit: Although.... after looking around a bit I realize this might not be the best idea to test for hydrazine....

Alkali metal hydrazindes, formed by dissolving the metals into hydrazine, are one of the few substances which behave as bases. These tend to be pyrophoric with air, and sodium hydrazide (NaN2H3), for example, explodes violently when heated above 100 C. As a result, alkali metals are incompatable with hydrazine. Alkali metals are common to reduce melting temperatures in some glasses-glass composites, for example.

From Here

Anhydrous hydrazine and sodium in ether react to form sodium hydrazine which explodes on contact in air....

From Here

So yeah, maybe it's a good thing I didn't get a yield.

[Edited on 9/27/2004 by BromicAcid]

BromicAcid - 9-10-2004 at 13:02

Immediately after the reaction mixture cooled after tying this synthesis I took it and cleaned it. But there was an adherent layer all in the reaction flask and going up into the still head that I could not clean off. I decided to soak it in concentrated HNO3 for a week or so.

Today when I pulled it out it was obvious what the gunk on the glassware was. Something had ate it somewhat terribly.

Hydrazine was made, not much but some. Maybe someone else with more patience then me can one day achieve a decent yield by this reaction.

Theoretic - 8-11-2004 at 08:06

Hydrazine can also be made by reducing nitramine. Nitramine is made by hydrolysis of dinitrourea, that being made by dehydration of urea dinitrate, that is made by reacting urea and nitric acid (ask your buddies at E&W :cool: ). The reducing agent could be sodium sulfide or something.

Mongo Blongo - 9-11-2004 at 09:29

I can't seem to find any information on Urea Dinitrate. Does it form with concentrated HNO3 rather than just Urea Nitrate at lower concentrations? I assume it can be precipitated with water right?

Theoretic - 9-11-2004 at 15:24

Have a look at E&W (if you don't know about it yet then google explosives & weapons), they have it in the High explosives section, the topic is called "Urea Nitrate and Nitrourea...again".

Mongo Blongo - 10-11-2004 at 11:58

Yea I have been there for a few years :). I read that whole thread again and still find no mention of urea DInitrate. Just urea nitrate, nitrourea and dinitrourea.
My goal IS dinitrourea as you can read in the "KETO-RDX and DIKETO-HMX" thread on the E&W forum where I have posted about a few experements.
So far I can only get to nitrourea and then it goes wrong for some reason. I am just looking for another way to achieve the dinitrourea synth.

Theoretic - 10-11-2004 at 13:16

I think urea dinitrate can be prepared in much the same way as urea nitrate, you just use twice as much nitric acid.
It's quite a straightforward procedure, essentially a formation of diuronium nitrate CO(NH3+)2(NO3)2

PainKilla - 12-11-2004 at 05:22

I will be attempting to perform the hydrazine synth using the NaOCl, urea and NaOH method. I am told that regular 6.25 is great for this because the reaction proceeds and is not too violent. However, if this does not work, would Ca(OCl)2 work? I can't figure out the equation for it. I haven't taken chemistry yet, all i know is from reading so pardon my dumbass :). Also, adding dilute h2so4 would form the sulphate without hazardous explosion hazards, yes? I will report fairly soon :p. Thanks.

Theoretic - 12-11-2004 at 07:46

A useful link

BromicAcid - 12-11-2004 at 10:07

Ca(OCl)2 is somewhat less soluble then NaOCl and I have never seen a preparation use Ca(OCl)2, it might be advantageous because the slow rate of reaction due to slow solvation, but the alkaline environment might just pacify the top layer. I've always wanted to try with Ca(OCl)2 because it is more readily available to me, however safety was not my concern, I planed to distill the hydrazine azeotrope directly from the reaction mixture which results in a higher yield. If you do experiment with Ca(OCl)2 be sure to try many different reaction contidtions and concentrations and addition methods and keep close track of them, you could conceivably start a new trend in hydrazine production.

And yes, addition of sulfuric acid will precipiate out your hydrazine sulfate which is of somewhat low solubiliy. If you did it with calcium hypochlorite though you would also precipitate out calcium sulfate and contaminate your product, but you could always recrystalize from warm or hot water.

Hydrazine KClO3

SAM4CH - 10-12-2004 at 14:10

Can I get hydrazine hydrate from KClO3 instead of NaOCl solution? Please more details!

KClO3 ?

MadHatter - 11-12-2004 at 00:31

I don't know what would happen by mixing urea and a chlorate. I would NEVER mix a
chlorate with ANY compound of AMMONIA ! This may produce the very dangerous and
unstable NH<sub>4</sub>ClO<sub>3</sub>.

the distillation

october-sky - 11-12-2004 at 05:47

What about the distillation, i have all the compounds (without the nitrogen) but i dont know how to distill the hydrazine...
i only know how to make a simple distillation...

chemoleo - 11-12-2004 at 12:50

No, you can't use KClO3, as much as you can't use KClO4, or KMnO4, or whatever. It's the -OCl you are after, because it reacts with the NH3 in such a manner that chloramine H2NCl is transiently produced.

As to distillation - if you read the thread carefully - it is commonly recommended to PRECIPITATE the hydrazine with H2SO4. No distillation involved. I wouldn't recommend it anyhow as distillation sometimes leads to explosions.

Comon guys, read the threads before you post obvious stuff!

[Edited on 11-12-2004 by chemoleo]

NaOCl from NaCl!

SAM4CH - 15-12-2004 at 06:10

Can I get pure NaOCl with electrolysis of NaCl as a method like preparing NaClO3?!

Please I need more detail for the cell current, running time, and others to improve my Chlorate cell to make pure hypochlorite.

BromicAcid - 15-12-2004 at 10:05

I think the key difference between a chlorate cell and a hypochlorite cell would be the cell temperature. Keep it low and hypochlorite would be favored. You might want to keep the current density a little lower too to prevent further oxidation of the hypochlorite.

BTW, I've tried many times to prepare hypochlorite insitu for reactions with electrolysis with little sucess, being that hydrazine is somewhat sensitive to impurites, making your own hypochlorite might be benifical, however you will have to use electrodes that will be both inert to the oxidation process of the electrolysis and that if they leave contaminates in the solution will not affect hydrazine production.

[Edited on 12/15/2004 by BromicAcid]

Marvin - 15-12-2004 at 13:42

Might I suggest setting up a chlorine cell, or even just a chemical chlorine generator and leading that into sodium hydroxide. There is little point in having the hypochlorite generated in the electrochemical cell itself. The chloride formed as well should not affect anything.

Rosco Bodine - 15-12-2004 at 19:45

Another option for you if you can't obtain 10% pool chlorinator sodium hypochlorite ,
but do have access to ordinary bleach concentration 5% or 6.25% sodium hypochlorite , and also have available calcium hypochlorite and washing soda , is to increase the concentration of the bleach you can get . The strategy would be to dissolve in the bleach sodium carbonate sufficient for producing the further amount of sodium hypochlorite needed to bring up the concentration to the desired level . Then to the stirred solution add calcium hypochlorite sufficient to react with the sodium carbonate , and filter out the calcium carbonate byproduct to obtain your more
concentrated sodium hypochlorite . A slight excess of theory of sodium carbonate should be used to insure that all the calcium ion is precipitated as carbonate . Or a slight added amount of baking soda could be used to finish the reaction and cleanup any remaining calcium which may persist in solution .
Washing soda as obtained from the grocery store is a partial hydrate so you should perform a weight loss on heating and drying test to determine the actual content of Na2CO3 . For example I baked 500 grams of washing soda for a couple of days and it lost 83 grams of water .
You could test twenty grams of your calcium hypochlorite separately to see just what amount of your washing soda is required to precipitate all the calcium , until futher additions of the sodium carbonate solution cause no clouding when mixing , to arrive at an "equivalence" ratio for the two reactants . You can then use that figure and do the math to know how much of each will be reacted in solution with your bleach to increase its concentration to whatever you desire , but 10 or maybe 12% at most is all you really need for the hydrazine from urea and sodium hypochlorite process . I have not actually tried this proposed method , but it would seem to be more direct than producing gaseous chlorine and reacting it with sodium hydroxide . One thing you must do is to keep the temperature low as possible and not permit any warming of the reaction mixture to occur for any extended time of even a few minutes or the hypochlorite will be destroyed . Very moderate and preferably cool temperatures are favored , if the reaction will proceed in the cold , even better .

tokat - 18-12-2004 at 23:13

and any execcess of water the thing will ignite, Use small amounts of compounds.

My hydrazine attempt using Mr Anonymous' method

garage chemist - 21-4-2005 at 13:15

About two years ago I made hydrazine sulfate with Mr. Anonymus' method. From 475ml of 10% NaOCl I got 15g HS. It was a terrible mess, the amount of foam was about ten times the initial volume of the solution and the reaction mixture poured out of the 2 L Erlenmeyer over my bare hands on the table and on the floor (I was really careless back then). I hope I didn't suffer any nerve damage from the hydrazine. Had some fun when I searched on the internet for the dangers of hydrazine (of course after getting it on my hands).

But once I had the HS, the rest of the synthesis was fun and a piece of cake.
To produce a maximum concentration hydrazine hydrate, I mixed the dry HS with an excess of NaOH and ground it in a mortar and pestle and distilled it. I got several ml of slightly fuming liquid.
The glass round bottom distilling flask was heavily attacked, though. This was not ideal.

I didn't like the freebasing method into methanol, and I didn't have any methanol so I chose to distill it.

I reacted it with isopropyl nitrite in ethanolic NaOH, it became quite warm and the IPN boiled away! Be sure to use an excess of IPN and cool it in an ice bath once the reaction has started.
(IPN was extremely easy to produce by slowly dripping HCl into a stirred solution of IPA and NaNO2 in water at 0°C, the IPN floats to the top and is separated and washed with water and bicarb soln. I used molar amounts, the IPN yield was high.)

The final yield was 3g of white fluffy stuff that deflagrated when heated and left small metallic globules which burned with a yellow flame.

E. b. C.: title

[Edited on 1-5-2005 by chemoleo]

Rosco Bodine - 21-4-2005 at 16:24

Originally posted by garage chemist
About two years ago I made hydrazine sulfate with Mr. Anonymus' method. From 475ml of 10% NaOCl I got 15g HS. It was a terrible mess, the amount of foam was about ten times the initial volume of the solution

You had to be doing something differently from the way the writeup described , because a volume increase of triple is
about average , sometimes a bit more but
quadruple would be the maximum and very unusual .....never ten times , unless
a spill of hydrazine just has a way of looking like a lot bigger mess than it is ;)

and the reaction mixture poured out of the 2 L Erlenmeyer over my bare hands on the table and on the floor (I was really careless back then). I hope I didn't suffer any nerve damage from the hydrazine. Had some fun when I searched on the internet for the dangers of hydrazine (of course after getting it on my hands).
No biggie for the dilute solution . HS is even taken medicinally as an experimental cancer drug . Hydrazine is destroyed so rapidly
by the air , you could come back to the area in a few hours , and there wouldn't be any of it left . No decontamination required , oxygen took care of the problem .

But once I had the HS, the rest of the synthesis was fun and a piece of cake.
To produce a maximum concentration hydrazine hydrate, I mixed the dry HS with an excess of NaOH and ground it in a mortar and pestle and distilled it. I got several ml of slightly fuming liquid.
The glass round bottom distilling flask was heavily attacked, though. This was not ideal.

There is no need to distill . You have the freebase , so just extract it with alcohol
and then nitrosate it .

I didn't like the freebasing method into methanol, and I didn't have any methanol so I chose to distill it.

Microtek has reported at E&W that the hydrazine can be freebased into anhydrous isopropyl alcohol using a special technique and proportions .


I reacted it with isopropyl nitrite in ethanolic NaOH, it became quite warm and the IPN boiled away! Be sure to use an excess of IPN and cool it in an ice bath once the reaction has started.

Everything needs to be cold at the start and through to nearly the end . When your isopropyl nitrite contacts the ethyl alcohol , transesterfication produced some ethyl nitrite , which is more volatile
and escapes more easily from a mixture that is not cold .


(IPN was extremely easy to produce by slowly dripping HCl into a stirred solution of IPA and NaNO2 in water at 0°C, the IPN floats to the top and is separated and washed with water and bicarb soln. I used molar amounts, the IPN yield was high.)

A slight excess of theory of NaNO2 to HCl should be used and a slight excess even further of isopropanol . This will result in a more stable product and even higher yield .

In my experience with making IN different ways , it is preferable to premix your ice cold alcohol and ice cold HCl , and to chill a concentrated solution of NaNO2 separately . Put the additional water which will be needed in the completed reaction mixture to insure the solution of the NaCl byproduct , as ice cubes in the empty reaction flask , and chill down everything in the freezer . When you remove the flask of ice cubes and pour over the cubes the very cold brine of NaNO2 , the mixture self cools from the salting effect , as the HCl and alcohol is
run in rapidly through a tube extending to the bottom of the stirred mixture . You can make 200 grams of IN in 15 minutes by this method , and it separates pure in a floating layer with the nearly melted ice ,
which rinses it as it melts , right where it is floating . If you use some solid ice cubes in the reaction mixture , they float in a layer with the precipitated IN and rinse the IN in situ . All you need to do is separate the product with the nearly melted cubes by decantation into a funnel , drain off the lower aqueous layer , and dispense the IN into a cold bottle having a teflon gasket , or a glass stoppered bottle . IN attacks anything else . The IN can be stored in the freezer .

It is very possible that the IN does not even have to be made separately if the hydrazine has been freebased into anhydrous isopropanol . The isopropanol
solution of hydrazine hydrate is basified with NaOH in equimolar amount to the HS which was freebased , and then the basified solution is treated with N2O3 ,
forming the IN in situ as a transient species , or reacting directly with the hydrazine , in either case causing NaN3
to precipitate . Microtek reported success
producing IN by bubbling through cold isopropanol N2O3 evolving from a warmed mixture of ~60% HNO3 and starch , via a smooth and steady decomposition of the HNO3 . Producing and isolating the IN separately is possibly
not necessary , if you follow the logic ,
but simply bubble the N2O3 directly into
the basified hydrazine extract in isopropanol . No one has tested this idea
so far as I know .

The final yield was 3g of white fluffy stuff that deflagrated when heated and left small metallic globules which burned with a yellow flame.

You got some NaN3 , but wouldn't you like to have more grams for your effort ?

More thoughts about the HS procedure .....

The formation of hydrazine from urea is
definitely not without an indication that
some small reaction is occurring ;)
It is almost unbelievable until you see
it yourself isn't it , just how much foam
can be afoot quickly as if it were a stampeding hydrazine " moose " :D

The first time I made hydrazine by the urea - pool chlorinator method , I had an overflowing mess too , so I kept working at refining the method until I had it down to an art . It very likely will make a mess if you start the reaction without allowing
it plenty of " room to grow " :D , and even
CYA with some overflow and return path like a really large funnel fitted snugly
into the neck of the flask . But if you do follow the directions given at great length
in the " improved " synthesis writeup ,
good results and yields will be predictable . Once you see it does work
in the way described , it will be a breeze to make up HS , even in large batches .

I have kilos of hydrazine sulfate as the product of many experiments studying the reaction , and it really is a simple synthesis after learning how the reaction behaves and how essential it is to allow it plenty of space .

When I make the HS , I use a half gallon
of hypochlorite in a batch run in a 4 liter
Erlenmeyer with a like capacity overflow funnel . I have a 12 liter flat bottom Florence so I could probably run a batch
using a full gallon of hypochlorite at one time and keep everything tidy by fitting
a five gallon plastic bucket with a rubber bushing and using it atop the flask as an overflow and return , in case the foam became excessive . Not exactly microscale , but saves work over several smaller batches . However unless you have some insane need for huge quantities of HS , a 4 liter erlenmeyer batch size is plenty for most folks .
Even with your 2 liter erlenmeyer you can make whole lot more in a controlled fashion than 15 grams . You should have gotten at least 55 grams . Keep the batches small until you know the reaction well , and then you can maximize the yield later . The reaction is sensitive to small variables which you will see if you do the synthesis a few times .

[Edited on 22-4-2005 by Rosco Bodine]

garage chemist - 23-4-2005 at 06:26

Thanks for the interesting info!

I just stuck 475ml of NaOCl in the freezer and dissolved 0,7g gelatin and 45,5g urea in 50ml water. Tomorrow I will try the HS synthesis another time.

I will dissolve the 64g NaOH in 60ml water and add it to the ice-cold NaOCl in liquid form to reduce heat evolution. I hope this will allow me to add all the NaOH at once.

A question: how did you stir the mixture? I only have a 5cm magnetic stirring rod and the 475ml NaOCl only stands about 3-4cm high in the 2L Erlenmeyer. So only the inner portion of the liquid is stirred, and the outer part moves rather slow. Should I still add the urea/gelatin solution all at once into the "vortex"?

EDIT: This was written by microtek in the E&W (I'm not registered there, got acess to the forum by using the google cache) :

"For sodium azide production I use a method which gives hydrazine in alcohol ( from hydrazine sulfate ) with little or no water:

- 1 mol dry HS is placed in a flask along with a suitable amount of anhydrous isopropanol.
- 1 mol of NaOH pellets are added and the contents are triturated with a glass rod until they begin to react, forming a slurry of hydrazine hydrate and NaHSO4. This doesn't mix with the iPrOH, but forms a sticky goo on the glass.
- Another 1 mol NaOH is added which converts the NaHSO4 to Na2SO4 ( and forms a mol of water ) which separates cleanly as a white powder.
I would think that the Na2SO4 is a good enough dessicant to dry out the solution which can then be decanted.
- Another extraction or two with dry iPrOH recovers most of the hydrazine.

This alcoholic solution of hydrazine works well for producing sodium azide with isopropyl nitrite."

Is this the method you're referring to, Rosco?
How much isopropanol should be added?

[Edited on 23-4-2005 by garage chemist]

For best results follow Mr. A's method *exactly*

Rosco Bodine - 23-4-2005 at 07:49

Because that method has been carefully optimized already , and any thing you do to change things even in a small way will
cause a disproportionate negative effect .
Do the procedure the way it was described *first* to confirm it works , and then change things to prove the * small * changes adversely affect the reaction in much * larger * ways you wouldn't have expected :D Been there , done that .

Originally posted by garage chemist
Thanks for the interesting info!

I just stuck 475ml of NaOCl in the freezer and dissolved 0,7g gelatin and 45,5g urea in 50ml water. Tomorrow I will try the HS synthesis another time.

I will dissolve the 64g NaOH in 60ml water and add it to the ice-cold NaOCl in liquid form to reduce heat evolution. I hope this will allow me to add all the NaOH at once.

Hmmmm , that is counterintuitive because
that will *increase * the exotherm by heat of dilution at the very moment the reaction is initiating , and you will likely produce a geyser , instead of a slowly rising foam . Again , do it first the way Mr. A has described . The thermodynamics are already worked out for controlling the reaction , and they are as described .


A question: how did you stir the mixture? I only have a 5cm magnetic stirring rod and the 475ml NaOCl only stands about 3-4cm high in the 2L Erlenmeyer. So only the inner portion of the liquid is stirred, and the outer part moves rather slow. Should I still add the urea/gelatin solution all at once into the "vortex"?

You quickly pour in a steady stream the urea and gelatine into the well stirred cold basified hypochlorite . Just pour it in
quickly , " in a lump " as the expression
goes . You should have one to two minutes before the foaming begins ,
so you have plenty of time to mix the two liquids if you pour the urea-gelatine solution in without any delay , the idea being to get it all in and mixed before the
foaming will complicate adding anything .

You should really have a 3" stirbar , even in a two liter flask , the octagon sort 3" X 1/2" should do okay . And yes , at first you have to reduce the stirring speed to keep the shallow pool from splashing and just pull as good a " swirling " shallow vortex as can be managed . The speed can be increased later as the foam begins to rise so that the stirbar is actually churning through and tearing apart the foam and helping to break apart the bubbles so that it returns to a liquid state and very slowly settles out again as a lower layer , over the course of the reaction , which may take a half hour or more . Everything about what happens has been described in Mr. A's very detailed description . Trust the information as a lot of work and many experiments were done to pin down all those important details .

The sequence of preparation of the reagents is important , the temperatures
and how the prepared reagents are added together is also important .
The reaction * is * sensitive to small variables and they have all been covered
in the described procedure , so don't deviate one iota in the belief that it probably won't make any difference , because just the opposite is probably true instead :D

And please , for anybody who makes arbitrary changes in one of Mr. A's or
Rosco's described experiments and then gets a different result , understand that what detail was thought not to be consequential , was indeed consequential . For every process where it is described what works , there are unwritten volumes about variations on the method which don't work , and when something is changed about a procedure ,
it usually does matter , even if the reason is not obvious why that change should have any bearing on the outcome . The troubleshooting of the variables has usually already been done for those procedures where exacting methods have been detailed . So roll that ready made wheel first and see how it goes , before
reinventing it in any detail and then saying it isn't round :D

garage chemist - 23-4-2005 at 07:55

Hmm, I understand that the method is optimized, but I can't imagine that the way the NaOH is dissolved in the NaOCl is of such importance. The only important thing is to keep it cool, right? Also, there is no gas evolution in this step. Urea and gelatin are added much later!
Perhaps this is a misunderstanding?

The_Davster - 23-4-2005 at 08:40

I must say, Rosco your HS synthesis does work very well. I attempted it for the first time last night and it went almost exactly as desribed. I did it on 1/32 the scale so yields will be lower from mechanical losses, but the synthesis went smoothly. Even the foam heights you described were the same, which while that will not surprise you, it did surprise me.

Rosco Bodine - 23-4-2005 at 08:53

Originally posted by garage chemist
Hmm, I understand that the method is optimized,

Correct . Don't change things in an optimized method .


but I can't imagine that the way the NaOH is dissolved in the NaOCl is of such importance.

Because the heat of dilution is huge and
it will cause too much heat to add
it all at once , both for the stability of the hypochlorite and for the controlability of
the reaction producing hydrazine . You are
also adding water to the system by predissolving the NaOH . This alone will
diminish the yield and is compounded by
the other problems which will be caused .


The only important thing is to keep it cool, right? Also, there is no gas evolution in this step. Urea and gelatin are added much later!
Perhaps this is a misunderstanding?

You are trying to apply linear two dimensional thinking to a much more complex scenario than lends itself to such analysis . You can trust me on that or you
can prove it for yourself with failed syntheses by testing your theories in
the flask . Lets not debate the theory .
Better to talk about the comparisons of
experimental results later , and what may be concluded from what actually occurs .
Anyone is welcome to make whatever improvements on the methods are valid ,
and describe their improved yields and
all of the details about how this was achieved . But that is for later , after the
proof is established by the results of their
experiments , not by estimates in advance
of what results are unknowns until such
experiments are completed .

rogue chemist :

Thanks for the confirmation .

I know my own work pretty well on procedures and methods that are
either completely my invention or
are variations and /or improvements
upon already known methods . And
I'm never afraid to say " I don't know "
but this isn't one of those times :D

[Edited on 23-4-2005 by Rosco Bodine]

garage chemist - 23-4-2005 at 09:02

I just find it easier to add a NaOH solution to the NaOCl (maybe even drop- wise) than to add solid NaOH, which dissolves very slowly and the solution needs to be chilled between the additions.
But if you say that it is so important to add the NaOH in solid form, I will do so.

In my synthesis (2 years ago), the foam was produced INSTANTLY when I added the urea/gelatin. Maybe the NaOCl was too warm? The ice on the outside of the erlenmeyer had just melted.
Does adding the urea/gelatin to the NaOCl at sub- zero temperatures greatly diminish the yield?

I will attach a 3-inch glass rod to my small stirring rod to make a big one. I hope this helps.

Rosco Bodine - 23-4-2005 at 09:58

Originally posted by garage chemist
I just find it easier to add a NaOH solution to the NaOCl (maybe even drop- wise) than to add solid NaOH, which dissolves very slowly and the solution needs to be chilled between the additions.
But if you say that it is so important to add the NaOH in solid form, I will do so.

Yes it is easier , and I have done it that way too , once . It decreased the yield ,
and the hypochlorite still required being
prechilled , because the heat of dilution is still huge , even if the NaOH is initially dissolved in a minimal amount of water .
The NaOH dissolves easily in the well stirred very cold hypochlorite . So there is
no reason not to do it the way which gives
the best result .

In my synthesis (2 years ago), the foam was produced INSTANTLY when I added the urea/gelatin. Maybe the NaOCl was too warm?

Too warm or you didn't use distilled water
for the urea-gelatine , or your stirring was inadequate and/or the addition of the urea-gelatine was done too gradually .
Usually there is at least 30 seconds before
the foaming begins to rise . The mixture
clouds almost immediately and lightens in color from a gazillion microscopic bubbles ,
self-heating accelerates and then a surge
of rising foam steadily rises .

The ice on the outside of the erlenmeyer had just melted.
Does adding the urea/gelatin to the NaOCl at sub- zero temperatures greatly diminish the yield?

It slows the reaction to initiate at below about 3 to 5 C , slows the induction and extends the total time , decreases the yield slightly , and can actually increase
the total volume of foam produced . The
thermodynamic curve the reaction follows
is something of a tradeoff between reaction rate , volume of foaming , and yield , which has a sweet spot a few minutes after the reaction is begun where
the increasing temperature actually causes the foaming to dissipate more quickly . When you start the reaction too cold , the gentle exotherm of the self-reaction doesn't allow the mixture to heat up on its own to reach that desirable temperature , where the foam is destabilized and begins to disintegrate on its own . This extends the overall time for the reaction and the volume which the foam expands is increased , above what total volume would be if the initial temperature was higher .
But if the reaction is started too warm , then the foam height increases too rapidly and again the total volume is increased .
So there is a " window " temperature range for the reaction start that is ideal for a particular batch size , and stirring efficiency , and it will differ for a different batch size because of the different radiational heat loss and latent heat characteristics for a different geometry
for that different sized reaction mass .


I will attach a 3-inch glass rod to my small stirring rod to make a big one. I hope this helps.

You could probably put your smaller stirbar inside of a hollow poly tube like
an old marker pen or similar into which it would fit snugly , and provide a bit of extra diameter and length . Size does matter :D although you will have to watch the speed to keep it from uncoupling in
the mixture . There's nothing in this reaction which should terribly attack the
poly , although I haven't tried this to make certain . Just make sure you keep the stirring going throughout the reaction . You will see the spiraling bands
in the foam indicating that the entire mass of foam is actually moving , especially just before it begins to disintegrate of its own accord as you are
ramping up the heating .

Something you really should do is provide some expansion room and return path for any potential overflow . If you dont have
a large funnel handy , then cut the bottom
out of a 3 liter PET soda bottle and invert it and secure the open neck of the soda bottle in the open neck of the flask with
a sleeve bushing cut from a large diameter
hose , or perhaps simply duct tape the necks of the two vessels together . Whatever overflow occurs will rise up into
the soda pop bottle and then flow right back down into the flask as the foam disintegrates with passing time . Sit a small soup bowl or cereal bowl or saucer
in the opening where you have removed the bottom of the bottle , after you pour
the urea-gelatine down through the opening and into the flask . Just dump and cover . You should be able to scale up at least 50% from the 475 ml , and perhaps 75% more without any overflow .

Whatever the scale you use , keep to exactly the described relative proportions
on every ingredient . The amount of gelatine will have significant bearing on
the reaction so weigh it carefully and make sure your glassware is very clean
and definitely use only distilled water .
Metal ions poison the reaction .

[Edited on 23-4-2005 by Rosco Bodine]

The_Davster - 23-4-2005 at 21:25

Well shit, I may have hydrazine poisoning. I have had a headace and had some nausea today, symptoms of hydrazine poisoning.:( As well I have felt weird today, not a way I have felt before. I wore a respirator the entire synthesis, except during filtration.

It could just be that I'm a bit of a hypochondriac, and this is just a placebo type effect. But I am still worried.


Rosco Bodine - 23-4-2005 at 22:34

It is unlikely for there to be any acute poisoning .

If you used the quantities of acid specified
then there is no free hydrazine present past the end of the first half of the H2SO4
addition . And if you were using the HCl for the first part of the neutralization , the vapors from the HCl neutralize most of any hydrazine vapors , and keep the vapors in the flask as a sort of smoky fog similar to what you see when ammonia fumes and HCl vapors contact . The synthesis is a very low emission , low fume method , so low that I don't even use a respirator or powered ventilation , but simply do the reaction in an open area like a patio .
All I have ever done is use ordinary caution and not keep my nose stuck in the middle of the reaction , set the controls and step back from it , sit in a chair maybe ten feet away and watch the reaction run , when it is attended at all .

You could have a chemical sensitivity / allergy sort of reaction , but it is extremely unlikely for you to have been
poisoned even if some of the mixture got onto your skin . The reaction mixture is
not volatile nor concentrated enough to
pose any serious danger , except to someone who may have a peculiar sensitivity like an allergy .

garage chemist - 24-4-2005 at 03:48

Is it possible to run the diazotation of hydrazine entirely in isopropanol? Then transesterification won't be a problem.
What is the solubility of NaOH in anhydrous isopropanol? I fear it might be too low...
But KOH dissolves much better in alcohols than NaOH, this might be an idea. This would yield potassium azide, of course.
KN3 gives metallic potassium on heating, so one might find it more useful than NaN3.

Rosco Bodine - 24-4-2005 at 06:10

Originally posted by garage chemist
Is it possible to run the diazotation of hydrazine entirely in isopropanol? Then transesterification won't be a problem.

I believe it is very likely possible to skip the synthesis of the isopropyl nitrite , and to simply bubble N2O3 through the cold , basified isopropanol extract of hydrazine ,
from which the NaN3 should precipitate directly in crystalline form .


What is the solubility of NaOH in anhydrous isopropanol? I fear it might be too low...
But KOH dissolves much better in alcohols than NaOH, this might be an idea. This would yield potassium azide, of course.
KN3 gives metallic potassium on heating, so one might find it more useful than NaN3.

NaOH has fair solubility in isopropanol , good enough for this reaction , but not nearly so good a solubility as in ethanol or methanol which is even better still .
Of course those alcohols give rise to the transesterfication problem . KOH would
work , but how much would be gained in
terms of the lesser solubility of the KN3 in
alcohol would possibly be offset by the
added water content , since the KOH will
always contain more residual moisture .

My experience with isopropanol is that you
have to keep any extract absolutely as water free as possible , especially with any solid solubles present also which would be more soluble in the moisture ,
because a phase separation is likely to
occur . The water content can separate
from the isopropanol carrying the water
soluble materials with it . To get an idea
of what I am describing , you can add ordinary NaCl to a bottle of 70% isopropanol and salt out nearly all of the water as a separated brine layer . So
it is clear enough that the extract in isopropanol must be nearly anhydrous at
the beginning , and kept that way by being a fairly dilute extract , in order for the solution to remain a single phase which precipitates a dry crystalline product of high water solubility such as NaN3 . It is a reaction condition where the product could easily separate as a
liquid brine , instead of an easily filterable solid , if the moisture content was only a little too much in the reaction . So the gain for losing the transesterfication problem using isopropanol is that the nitrosation doesn't need to be run in a low pressurized system , which simplifies the apparatus requirements . But the use of isopropanol creates a higher demand for a more anhydrous condition for the extract , and because of the reduced solubility of the reactants in the
isopropanol the bulk yield possible from a given volume of reaction mixture is much
lower for isopropanol . If the methanol
extract can be managed with a pressurized apparatus , the yields from a liter of the methanol based reaction system are 3 or 4 times what can be gotten from a liter of the isopropanol based reaction mixture , because of the
solubilities of the reactants . Worthwhile yields are gotten using either solvent or
even ethanol , but each has its advantages and disadvantages .

Microtek described his method using isopropanol and attached is a copy of
the text originally posted at E&W . I
would just link , but they are having
bandwidth problems so I hope this
attachment is acceptable .

An idea which I think would be worth experimenting , with regards to the direct
nitrosation of the basified alcohol extract
of hydrazine hydrate , is to use a methanol free denatured ethanol for the extraction , and do the nitrosation at
salted ice bath temperature . Using
ethanol would sort of be an intermediate
choice between methanol and isopropanol , which could have the advantages of each so long as the reaction mixture was kept very cold .

[Edited on 24-4-2005 by Rosco Bodine]

Attachment: Microtek NaN3 post from E&W.doc (27kB)
This file has been downloaded 1831 times

garage chemist - 24-4-2005 at 07:18

That's what I wanted to know, thanks!
So the solubility of NaOH in IPA is too low to be convenient, and using IPA brings the moisture sensitivity problem.

Mixing the IPA hydrazine extract with ethanol, dissolving the NaOH and adding the IPN (excess) under cooling will therefore be my method.

BTW, I just dissolved the NaOH in the NaOCl (as described in the preparation, in two steps) and quite a lot of white solid has precipitated! I assume this is NaCl. Adding 50ml distilled water didn't dissolve it all, and I didn't want to add more water.
The NaOCl was 2 years old, that may be the reason. It likely contains more NaCl than fresh NaOCl, and this now precipitated.
I hope this won't cause problems other than a slightly lower yield.

I've put my stirring magnet into an 8cm teflon hose which had to be slit lengthwise to allow the magnet to be inserted.
This allows for better stirring.

Rosco Bodine - 24-4-2005 at 08:32

Studying the tradeoffs involving volatility scenarios from transesterfication , solubilities for the reactants , ect . ....

The best method in my mind for simplifying
the apparatus requirements , while still
having a good product yield from a given volume of reaction mixture , would be to
use a methanol free denatured ethanol
for freebasing the hydrazine . And then
to either use isopropyl nitrite obtained by
either microteks method , or the more usual method using nitrite and HCl , to
nitrosate the basified hydrazine extract
in ethanol , with the reaction done in a
salted ice bath . Alternately , simply bubbling N2O3 into the basified ethanolic
extract of hydrazine , cooled in a salted ice bath might work . Using ethanol should be the compromise which is easier for the extraction not being as picky about the moisture , while providing better solubility and a more concentrated reaction mixture , and yet still having a not too low of a boiling point for the transesterfication product which would be ethyl nitrite , which boils at 17 C . Salted ice would keep the reaction cold enough if it was done slowly , to keep the nitrite in the reaction .

Methyl nitrite on the other hand , boils away at - 12 C , so it tries to escape even at very cold temperatures unless a bit of pressure is used to contain it , the loss of nitrite is probably 50% , which is unacceptable .

So a variation on microteks method , but using ethanol for the extraction could prove useful , but only for denatured alcohols where methanol is not the
denaturant . It is possible that some
other denaturants might be problematic ,
and this would have to be examined .
It is also possible that a methanol denatured ethanol could be stripped of its methanol in advance by bubbling through
it sufficient N2O3 to esterfy and volatilize
any methanol . Warming the alcohol would drive out any residual methyl nitrite
leaving essentially pure ethanol . This same methanol removal method could be useful for purifying methanol denatured alcohol for use in fulminate synthesis , or other uses where the methanol denaturant is undesirable , and ethanol
alone is what is required . Hmmmm , the
revenue folks are sure to love this news :D Transesterfication purified moonshine anybody ? Pass me the jug ;)

Your two year old NaOCl is not NaOCl anymore . The storage life for the 10% is only a matter of weeks at cool temperatures . At subfreezing temps it
could be stored longer , but normal shelf life at ordinary temperature is very short .

[Edited on 24-4-2005 by Rosco Bodine]

garage chemist - 24-4-2005 at 08:49

Well if I add acetone to my 2 years old NaOCl, it still gets hot and starts boiling and splattering. 5% NaOCl only gets moderately warm, not warm enough to boil off the chloroform.

Another question: after the reaction, do I need to add exactly 164ml of HCl or can I just add HCl until the gas evolution stops?

Rosco Bodine - 24-4-2005 at 09:23

Keep to the stated proportions regardless
of the physical observations . There is some overlap for the -Cl and -SO4 neutralizations so that the CO2 evolution is not a precise indication of where the neutralization actually is in its progress .
I also believe the hydrazine is already largely tied up as a carbonate in the mixture , or as a complex with other carbonates . So you will get some CO2 evolution beyond the point where you would expect it to stop , even into the addition of the H2SO4 . But the ratios
have been worked out to be optimum for
the precipitation of the HS at the end ,
in pure form . So use the quantities given .

Do a molar quantities analysis , neutralization equivalents and some extra
for operating margins , and you will see the arithmetic is already worked out correctly . Quantity variations were tried for experimental confirmation that the math was applicable , which it was confirmed to be . I covered all the bases
on this HS synthesis . You can confirm that easily enough . It's like you are asking " are you sure " ....yeah I'm sure :D

garage chemist - 24-4-2005 at 11:08

OK, I just performed the synthesis.
I tested my NaOCl with acetone another time, and it became hot and a good amount of chloroform appeared in the flask. But it didn't start boiling anymore (the last test was a few months ago). I estimate my NaOCl at around 7%.

I poured the urea/gelatin into the basified NaOCl at around 4°C. After 15sec. it became milky white and a bit of foam appeared.
But after 5min. only a small amount of foam had been produced, and the solution still was rather cool. I heated it, and this started the foam production. The solution roughly doubled its volume (the foam got to the 1000ml mark) and became hot from the reaction (it could just be touched). The generation of only such a small amount of foam was the first remarkable thing. After 10min I heated it further, and the foam subsided completely. The solution was dark yellow/green. On further heating the color slowly became lighter. I heated it to about 80°C in 30min, then it was nearly colorless.
I cooled it to about 30°C, and added 217ml of 25% HCl. White fumes were produced in the reaction vessel, and lots of CO2 was produced. Then 50ml conc. H2SO4 diluted with the same amount of water were added. There was still lots of effervescence until nearly the end of the addition.
Then a white precipitate rapidly appeared.
The solution was still quite hot! This was the second remarkable difference to my first try 2 years ago.
On cooling, a lot more precipitate formed. It is definately more than 15g! I'm expecting at least 25g.
The solution is in the cooling bath at the moment. Tomorrow I will further cool it down with ice to precipitate all the HS.

Rosco Bodine - 24-4-2005 at 14:03

Deep Orange is the color you should see when the foam is subsiding , almost the color of strong tea , but more orange than
red brown . That color gradually fades to
a very light ale color , almost clear at near the endpoint .

The quality and freshness of your NaOCl is really what is the single most important
factor affecting yields if you follow the rest of the procedure exactly . That is the one factor over which there is little control , except that you can generally expect the fresh stock arriving at stores
from the spring through to the end of summer during the swimming pool maintenance season to be good quality .
If you plan on storing the 10% for any time , keep it in the refrigerator or freezer . Outdoors in warm summer temperatures it will be useless in about
2 weeks . That temperature sensitivity
is exactly why the NaOCl must be freezing cold when you dissolve in it the NaOH ,
and do it gradually so as to produce no
undue warming of the solution , because the heating decomposes it on the spot ,
before it has a chance to participate in the synthesis of the hydrazine by doing its
decomposition in reaction with the urea .
The NaOCl is also pH sensitive , and it only can exist in alkaline condition . So
NaOH doesn't hurt it , actually even stabilizes it , but pH in the other direction
can immediately cause complete decomposition of the NaOCl , a situation
aggravated by warm temperatures .
Your old stock will gradually decompose to
a mixed solution of NaCl and NaClO3 ,
the ratio depending upon the pH and temperature . So use only the fresh stock for the HS synthesis if you can obtain it .

With some fresh NaOCl , and your 2 liter flask , if you did make an overflow funnel like I suggested using a 3 liter pop bottle ,
you could probably scale up to a .35 amount of the materials used in Mr. A's
improved method , if your stirrer is doing the job . Work with 665 ml of the NaOCl
and you should get about 80 grams of HS for your trouble . You can probably start that reaction a bit warmer if it isn't taking off on its own exotherm and maintaining the reaction . Anywhere around 6 - 8 C should work fine . The fresh hypochlorite
is going to be more reactive , and the larger reaction mass should keep going from its own exotherm without any help in the initial stage of the reaction .

[Edited on 24-4-2005 by Rosco Bodine]

Mumbles - 25-4-2005 at 17:46

I had a very odd happening when I made my first, and only batch of Hydrazine sulfate. I used the proceedure on Rhodium's then existant site. I scaled everything down to 1/10 the values and modified the Hypochlorite level to account for my differing concentration.

There was absolutly no color change like described or foaming. It changed colors almost unnoticably. There was a yellowish-green color at the beginning, no doubt from the bleach. It then changed slightly darker orange, and more of a yellowish near the end.

This experiment took place a while ago, but I have been vacant from this fine forum for a bit of time. Does anyone with some experience in the proceedure know what was going on? Would the lower concentration or smaller ratio cause nearly no color to be imparted into the reaction, and there to be a complete lack of foaming?

Polverone - 25-4-2005 at 18:41

If you used the procedure posted on Rhodium's, be aware that it was transcribed from this site. So basically read all that Mr. Anonymous / Rosco have to say about the subject, then ask Rosco or the others who have followed his methods if you still have questions. Mr. Anonymous's messages were posted under my username, so take that into account when searching.

Rosco Bodine - 25-4-2005 at 19:00

Are you talking about the urea and hypochlorite reaction , having no foam ?
That would seem impossible .

It would make sense for there to be less
instense color tint developed in a reaction
mixture which is diluted to about half the usual content of reactive components .

I haven't tried the reaction using ordinary
household bleach . But using the 10% pool chlorinator , the color reactions have been consistent , always the orange color
appearing , without exception , and always there has been plenty of foaming .

Mumbles - 25-4-2005 at 19:18

Yes, it was urea and hypochlorite. I have no idea what happened. I heated it exactly as the proceedure called for. It was very similar to your proceedure actually. It came as a total suprise when there was not one bubble of foam produced.

I had expected for the colors to be somewhat less intense, but this was way too light from what I was thinking. The pictures I've seen of the reaction and what I got were totally different. This was about the intensity of color as to what a good titration end should be, just barely visible. I had to hold it up to the light to see any colors at all.

I got about 5-7% lower yields than what was described, which is expected from the lower concentration and smaller reaction size. This is the most suprising fact to me. It didn't behave at all as was expected, yet I still got a normal yield. The product is primarily Hydrazine as well. There was definatly a period of temperature that was cooled after filtering the HS, before the sodium sulfate precipitated.

I will make another attempt soon using my same reagents and proceedure and see if there is normal results, or if I recieve the same ones I had last time. I haven't a clue as to what happened.

Rosco Bodine - 25-4-2005 at 19:44

That is very interesting . I would still recommend trying the proportions and the
use of HCl for the preliminary neutralization , as described in the improved method , as it should result in
better economy and improved crystals .

It could be that if you didn't increase the gelatine to account for the increased volume of the reaction mixture , that the viscosity of the mixture was low enough so that the foam would not occur , and that the CO2 simply effervesced cleanly
without producing a foam .

I have actually wondered if there is not some sort of dispersion agent which could
be added to the more concentrated mixture to destabilize the foam and reduce the nuisance it presents . I tried reducing the amount of gelatine which helps , but it also reduces the yields .
Perhaps another different chelation reagent could be substituted for gelatine
which would eliminate the foaming completely .

garage chemist - 26-4-2005 at 06:22

My problem is that NaOCl pool chlorinator is not available OTC where I live. They only have TCCA and H2O2 based stuff.
2 years ago I was able to buy 10 Liters of industrial chlorine bleach, which guaranteed to contain at least 14% NaOCl at the moment of delivery. (The company doesn't sell to individuals anymore now).
Stupidly, I didn't adjust the ratios. I think that most of my produced hydrazine was oxidised by the excess of NaOCl. This would explain all the foaming (nitrogen!) and the very low yield.
Now, the solution is old and weak.

My only source of NaOCl is the 5% NaOCl OTC bleach, which is actually rather cheap.
I will conduct experiments for hydrazine production from this.
I will use the same procedure as with the 10% NaOCl, I will only use double the amount of 5% NaOCl bleach. Everything else will be the same (amount of gelatin, amount of urea etc...).

A 5% NaOCl solution is actually nearly indefinitely storage stable, the 10% pool chlorinator will eventually reach this concentration and then remain at it, as I've read at several places.

BTW, silicone oil is a very effective foam suppressant. One drop is enough. Octanol also works.
But I'm not sure if this will help with the very rigid gelatin foam.

Rosco Bodine - 26-4-2005 at 08:34

There is an " extra strength " household
bleach usually labeled as " Ultra " bleach
and it is about 6 % while the regular bleach is usually 5.25 % NaOCl . The reaction for hydrazine should still work to some extent and you can experiment to see what sort of adjustments may benefit the reaction with the lower concentration bleach . The disadvantage will be the lower amount of hydrazine sulfate which is possible to get from a given volume of the reaction mixture . If the yields are similar percentages on a molar basis , then the economy of the reaction will be
unaffected , except for the increased labor and need for larger reaction vessels . Of course , if the foaming problem is gone then the volume requirement for the reaction vessels will
not be reduced because of the larger volume required for the more dilute solutions for a similar batch size . And after the neutralizing acids are taken into account , the volume efficency will be reduced . But if the process is still workable even with ordinary bleach , that is certainly a valuable development for
being able to substitute a more readily available form of NaOCl .

It may be possible to make NaOCl of high concentration by reacting calcium hypochlorite pool shock with sodium carbonate washing soda , filtering out the precipitated calcium carbonate from the calcium compounds , leaving a solution of NaOCl . It may be possible to boost the NaOCl content of ordinary bleach by this method . One would need to do a weight loss on heating test for the washing soda to determine its state of hydration . The
Arm and Hammer brand which I tested is typically 83.4 % Na2CO3 and the remaining weight is water of crystallization
in the form it is found straight from the box . Other brands or lots may vary .

So this is a possible adjustment method which may be useful if it is needed .
However , if useful yields are still obtainable from ordinary strength household bleach , then it probably easier to just work with what is available and
simply use larger batches since the process is very scalable up or down .

I have tossed around the idea that in an insulated container such as a large plastic lined picnic cooler , that a certain volume
and geometry reaction mass could probably achieve the conservation of its own heat of reaction sufficiently for the
exotherm to be sufficient alone to carry the reaction through to completion without any supplemental heating being required . If this is true , it may be easier
to make HS in kilogram quantities in one batch , using several gallons of bleach at
one time in a large enough scale batch to
achieve that conservation of the exotherm . All that would be required then is to let the reaction run on its own exotherm to completion , and then neutralize the mixture . On an industrial scale there seems no reason why this would not work , and down to a certain
limiting scale of perhaps three to five gallons of mixture , the same method may be adaptable to " small scale " method which is possible in " picnicware " sized
insulated containers :D

garage chemist - 27-4-2005 at 02:08

I weighed my yield: 29 grams!
And this from 2 years old NaOCl.
From the theoretic yield for fresh NaOCl, I calculated my old NaOCl to be at around 5,3%.
Experiments with 5% OTC bleach will be conducted soonish.

Calcium hypochlorite isn't available where I live either (I live in germany).

I could theoretically make my own NaOCl by reacting chlorine (from TCCA) with NaOH, by this way, I could make up to 25% NaOCl which should give impressive yields. But that would be a lot of hassle.

A question: as I understand it, the freebasing gives anhydrous hydrazine in isopropanol.

From the reaction

N2H6SO4 + 2 NaOH -----> N2H4 + Na2SO4 + 2 H2O

we see that two moles of water are produced per mole of hydrazine.
But the formed sodium sulfate is a drying agent and binds water according to this equation:
Na2SO4 + 5 H2O -----> Na2SO4* 5 H2O

So the isopropanolic hydrazine solution should be anhydrous, shouldn't it?

[Edited on 27-4-2005 by garage chemist]

chloric1 - 27-4-2005 at 05:00

Experiments with 5% OTC bleach will be conducted soonish.

A question: as I understand it, the freebasing gives anhydrous hydrazine in isopropanol.

From the reaction

N2H6SO4 + 2 NaOH -----> N2H4 + Na2SO4 + 2 H2O

we see that two moles of water are produced per mole of hydrazine.
But the formed sodium sulfate is a drying agent and binds water according to this equation:
Na2SO4 + 5 H2O -----> Na2SO4* 5 H2O

So the isopropanolic hydrazine solution should be anhydrous, shouldn't it?

[Edited on 27-4-2005 by garage chemist]

I would bet you would have the hydrazine hydrate in the isopropanol BUT they do use alkali to dehydrate hydrazine hydrate. Most reactions would not require anhydrous hydrazine anyway so if you wanted the hydrate I guess adding a little water to the extract and removing an isopropanol azeotrope under reduced pressure would be satisfactory.

P.S was not aware soonish was a word. I like it!:D

[Edited on 4/27/2005 by chloric1]

usual behavior of hydrated salt does not apply

Rosco Bodine - 27-4-2005 at 07:39

It would be nice if the alcohol extract of hydrazine hydrate was dried in situ by the
sodium sulfate , but the behavior of the
sodium sulfate with regards to hydrate formation is modified under these conditions . The temperature at which any hydration would occur is lowered perhaps 30% , and the formation of the hydrated forms of sodium sulfate is undesirable since the water of crystallization invariably carries hydrazine hydrate along with it , resulting in loss of hydrazine occluded in the crystals . The bottom line is the hydrazine is more hydrophilic and hygroscopic than is the sodium sulfate which would be used to tie up the water , and the hydrazine wins the tug of war for the water . So it is best to simply decant the alcholic extract of hydrazine hydrate , actually along with the additional mole of water ( it is probably hydrazine dihydrate ) , leaving the anhydrous sodium sulfate , operating at a temperature which precludes any hydrated salt formation . Better to get some water along with the hydrazine , than to lose hydrazine along with the water tied up by hydrated salt formation .

Freebasing hydrazine into alcohol is something of an artful manipulation of solids and working with the solubilities of the components while trying to eliminate completely as nearly as possible , having to actually add any free water to keep the mixture a paste which can be stirred .

Understand that there is a half-neutralization point for hydrazine sulfate , where a dihydrazine sulfate is formed , which has ten times the solubility in H2O
at 60 C , as does monohydrazine sulfate from which it derives . At 60 C hydrazine sulfate is soluble to an extent of 8.3 grams per 100 grams of solution , while the dihydrazine sulfate is soluble 84.7 grams per 100 grams of solution . So the mixture becomes thinner and more fluid to a point as if it was melting , and then gradually thickens and almost sets up solid as the freebasing is completed by the introduction of the last half of the NaOH .

A special strategy can be used to exploit the solubility change . You first introduce into the flask about half the hydrazine sulfate to be freebased , and then in portions add about a quarter of the total amount of NaOH , until a thin stirrable slurry is formed , which will be hot from the heat of reaction . Then in alternating portions to this stirrable mixture is added the remaining HS and NaOH , at a rate which maintains the mixture hot and keeps the mixture fluid enough to stir .
When all the solids have been added and mixed together well , and before the mixture has cooled to the point of setting up solid , the alcohol for the first extraction is added in a lump to the slurry , stirred well , and then decanted while still warm . This first extract contains most of the hydrazine hydrate ,
probably ninety percent or more of the
hydrazine , and what remains is only what is physically trapped in the wet slug of crystals and cannot be decanted . Subsequent extractions and decantations
of the alcohol simply reduce to a small remnant the amount of residue of hydrazine trapped in the wet crystals
which defy any filtration due to the air sensitivity of the hydrazine . The blanket
of alcohol fumes coming off the warm extracts affords some protection against air exposure during the brief times when the liquid extract is being decanted from one container to the next . In between such transfers , containers must be kept stoppered . This is quite essential for preventing loss of the air sensitive hydrazine . In the free form hydrazine is extremely reactive .

These are some pertinent references

GB900397 Freebasing hydrazine in ethanol

GB876038 Frebasing hydrazine in methanol

garage chemist - 27-4-2005 at 07:53

I just found out that the "5% bleach" is actually 2,8%. That is too low to be of use.
I will have to look around for a supplier for the 10% stuff. But I will first use up the 5,3% stuff, that will surely give me more than 100g of HS.

Thanks for the freebasing instructions, that was really helpful.
Ethanol could also be used for the extraction, I hope.

@ chloric1: I know my english is a bit funny, I sometimes use words which I read once, even if they are wrong or don't exist.

Rosco Bodine - 27-4-2005 at 08:11

My experience with the decomposition of NaOCl solutions is that once the decomposition starts , it is autocatalytic
and continues until the NaOCl content is zero , and only a solution of NaCl and NaClO3 is left as the product .

The decomposition doesn't " stabilize "
at some percentage like 5.25% and stop ,
but continues to complete decomposition .
It is pH related and temperature related .
Generally , the higher the initial concentration of NaOCl , the greater the instability and the shorter is the storage half-life , which is an inherited instability for the lowered concentration solution at any point in time .

The 10% bleach is used for many other purposes than as a liquid pool chlorinator .
It is also used as a sterilizer wash for recreational vehicle waste storage tanks
and piping . So you may be able to find it
in the camping supplies section in some large stores . Some people also use it in their pressure washers for killing mildew on decks and for pressure cleaning walls
and rain gutters on buildings , sidewalks ,
bathhouse / bathroom areas and I have seen it used as a washdown around dairies , delis and fish markets , butcher shops . Maybe a janitorial supply would have it . It seems like it should be a common item .

[Edited on 27-4-2005 by Rosco Bodine]

garage chemist - 27-4-2005 at 12:20

I calculated:

If 108g NaOH are dissolved in 796ml water and chlorine is bubbled through (while cooling the solution with ice) until no more is absorbed, one gets about a kilogram of a fresh 10% NaOCl solution.

Can somebody confirm my calculations?

Rosco Bodine - 27-4-2005 at 12:56

Stoichiometry gives me a headache :D

Sometimes a bit of background music helps .

Anyway , if you simply must resort to chlorination , it would probably be easier
to use regular bleach as the starting point , chill it feezing cold and go ahead and add NaOH in the quantity needed
for the added NaOCl , plus whatever extra
NaOH is to be added for the extra needed in the hydrazine synthesis . Rechill the solution , and chlorinate it with the vessel
sitting upon a scale until it has absorbed the necessary weight gain . The strongly
basified mixture will chlorinate more easily
and you will have one less step to do afterwards , plus the product will be more stable .

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garage chemist - 28-4-2005 at 00:42

I can as well start from plain NaOH solution than from the really weak NaOCl bleach.

What I don't understand is why you have to monitor the weight gain of the NaOH during chlorination! Why don't just bubble in chlorine until no more is absorbed?
EDIT: sorry, I missed the part where you said that the extra NaOH is used in the hydrazine synthesis. Then chlorination would have to be stopped at the right point.
But even weak NaOH vigorously absorbs chlorine, you can be sure that there will be no problems with unreacted chlorine even when you chlorinate until all NaOH has been converted.
My scale only goes up to 500g, so I can't make a kilo of NaOCl with weight monitoring.
Perhaps add some extra NaOH after chlorination to bind the free chlorine, which would otherwise decrease the pH and destabilize the NaOCl.

I can get TCCA cheap and in large amounts, so I can make all the chlorine I'd ever need.
Calcium hypochlorit isn't available, though.

BTW nice background music! :)

[Edited on 28-4-2005 by garage chemist]

Rosco Bodine - 28-4-2005 at 06:42

The factors which would have me suggest you use bleach as a starting point and strengthen that bleach are several .

The gallon of bleach is already at least halfway there to being what you want in terms of its NaOCl content , plus it also contains the distilled water you will also
require , along with the NaOH and Cl which would have to be purchased separately , and all three components then used to form what is the product
which is sold cheaply . Also , half of the
thermodynamic of a highly exothermic
process has already been managed ,
which is really the greatest technical difficulty . So it is for reasons of economics involving the cost of materials
and for technical thermal considerations ,
that it is better to strengthen regular bleach than to start from scratch . You can probably buy the already made bleach more cheaply than you could buy just the distilled water and sodium hydroxide required for making the same product .

Chlorinating to a specified weight gain of absorbed chlorine , with sufficient alkali in excess to maintain an endpoint at pH 11 or higher is the standard method for producing a stable product . If you chlorinate to a point of saturation , you will drop the pH and it will destabilize the solution producing chlorate instead of NaOCl . Also the water must be especially pure for stable NaOCl solutions ,
as even a few hundreths of one part per million of ions such as nickel , cobalt , copper and iron , cause catalytic decomposition .

The reaction for the formation of the Javelle water ( bleach ) , from chlorination of NaOH solution follows from the first reaction of Chlorine and water .

Cl2 + HOH -----> HCl + HClO

and in the presence of NaOH ,

( Cl2 + HOH ) + 2 NaOH ----> NaCl + NaOCl + 2 HOH

So , on a molar basis , ordinary bleach has
as much salt in solution as it has sodium hypochlorite .

By what method are you going to generate the Cl from the TCCA ?

[Edited on 28-4-2005 by Rosco Bodine]

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