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Author: Subject: Hydrazine
Al Koholic
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[*] posted on 23-11-2003 at 14:36
Hydrazine


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?
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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 23-11-2003 at 16:32
Qualitative & Quantitative Test for Hydrazine (Sulphate)?


Well, I just tried to make some hydrazine sulphate according to my post in the hydrazine oxidiser exploration thread ( http://www.sciencemadness.org/talk/viewthread.php?tid=376 ).
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]




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[*] posted on 24-11-2003 at 14:26


Silver mirror test and Greenwood and Earnshaw.
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[*] posted on 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.

unionised,
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.

chemleo,
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.
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[*] posted on 24-11-2003 at 23:48
Damn


Well Marvin...you 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]
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[*] posted on 26-11-2003 at 20:27
test


from THE CONDENSED CHEMICAL DICTIONARY:
copper dihydrazinium sulfate
  CuSO4(N2H4)2.H2SO4
 Properties: Bluish powder; mp above 300°C, starts
  to decompose at 140°C; very slightly solouble in
  watrer, 250 ppm @ 80°C.
 Hazard: Moderately toxic by ingestion or inhalation.
  Skin and eye irritant.
 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 ...




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[*] posted on 30-11-2003 at 03:19
Hydrazine


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
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smile.gif posted on 1-12-2003 at 17:52
Anhydrous Hydrazine synthesis, Unabridged for Newbies, Kewls and the other downtrodden misunderstoods...


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!




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[*] posted on 1-12-2003 at 18:20
Success with hydrazinium sulphate


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....




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[*] posted on 2-12-2003 at 20:16
A definite peculiarity...


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]
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[*] posted on 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]




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[*] posted on 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.
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[*] posted on 5-12-2003 at 17:11
More to the production of Hydrazine


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]




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[*] posted on 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!




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[*] posted on 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 :)
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[*] posted on 6-12-2003 at 08:24


Which is exactly why I didn't get any green gas because my solution had been freshly prepared!!!
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[*] posted on 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
Quote:

"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!




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[*] posted on 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.
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[*] posted on 13-12-2003 at 20:59
WOW


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)4 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




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[*] posted on 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: http://www.uspto.gov/index.html 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.




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[*] posted on 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.
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[*] posted on 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.
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