I tried to prepare anhyrous hydrazine from Urea with Iron powder as catalyst but I did not get any thing, I saw a white ppt. at my distilation tube
and ammonia through my distilation process, So please help me at this, and about Nickel powder I can't get it now.
ThanksThe_Davster - 16-7-2004 at 12:39
You could get Ni powder by reacting your iron powder(or an iron object) with a nickel salt.BromicAcid - 16-7-2004 at 13:57
What temperatures were employed. How fine was your iron powder? How about your urea, if it was not yet to molten temperature what form was it,
pellets? How long did you hold it at heat? What do you mean at your distillation tube? What did your setup look like, were you bubbling it into
something? Did you have the forethought to flush your reaction vessel with an inert gas?
All but the last question are incredibly important. In order to be any help I need significantly more detail. Note, hydrazine can explode from
ground glass joints, traces of other metals in your iron powder, or the sun hitting it just right. In addition it is a potent carcinogen and toxic in
other senses too, please wear proper protection.
BTW, aside from the two patents there is no proof this reaction works to any measurable extent, the patents give no yields so this may be a wild
goose chase.
[Edited on 7/16/2004 by BromicAcid]
More details
SAM4CH - 17-7-2004 at 00:59
I tried with 135-220 degree centigrate, and about my Iron powder I used rough powder (around 50 Mesh), My urea was melting at 135-138 degree and I get
a molten urea (I used technical grade) it take about 15 minutes in oil path, and about my distillation tube (I ment connecting tube joint between the
flask and graham condenser) and the white ppt. closed my tube, I try to distill hydrazine by distillation Column like used for Ammonia Distilling, I
do not want to make hydrazine sulfate I need anhydrous hydrazine directly from urea, I did not tried to use an inert gas at distillation.
Note: I saw an orange ppt. in my flask when I was stoping heating.BromicAcid - 17-7-2004 at 13:30
For the process with iron the patent calls for temperatures between 140 and 150 C so you should try to hold it here. At the higher temperatures the
urea may vaporize, may have been what clogged up your tube or possibly semicarbazide which can account for 20% of the urea consumed if this reaction
runs to completion. It is not a fast reaction it should be held at temperature for extended lengths of time. Also the reaction may be highly
dependent on surface area, 50 mesh iron might not cut it considering the lessened reactivity of iron in this reaction.
As for the orange precipitate, it could have been so many different things that I have no definite clue.
PS: How much iron powder did you use? It is customary to use a large excess, >50% by weight.
[Edited on 7/17/2004 by BromicAcid]thefips - 17-7-2004 at 15:16
Can someone please quote the synthesis?
As soon as I have more time,it would be a very interesting projekt.
Safty level
SAM4CH - 17-7-2004 at 17:59
Okay.. I used 50% by weight (60g urea + 30g Iron powder) So I like to ask what about the time needed to produce all anhydrous hydrazine (30g), what
about safty if I used glass distilling Column (with out an inert gas) and about toxic gas can I use charcoal to make high safty about Carbon monoxide
and Iron Carbonyl by passing gas around it, and what about all other catalyst can do this better than Iron or Nickel!.unionised - 18-7-2004 at 03:04
I must be missing something here. Why does the chemical industry carry on using the Raschig process if they could just piss on hot iron?
Is there any evidence that this method provides any useful yield of hydrazine?
It was very slow!
SAM4CH - 18-7-2004 at 06:20
I tried again with anhydrous hydrazine using iron powder (60g urea + 40g iron) and I saved temperature between 140-150 degree centigrate but the
decomposing of urea wsa very slow and I get a few drops in 3 hours (1drop every 20minutes), So please advise me to make N2H4 easly.FrankRizzo - 18-7-2004 at 13:52
Are you mixing the iron with the urea in your reaction vessel? If so, your iron catalyst is only being heated to the temperature of your decomposing
urea, and that is probably slowing the reaction down. IIRC, you are supposed to pass the decomposition products across hot iron particles (which are
heated externally).BromicAcid - 18-7-2004 at 17:03
Quote:
you are supposed to pass the decomposition products across hot iron particles
Where did you hear this? The patents state the combination of urea with the carbonyl forming metal and heating. Nickel supposedly yields hydrazine
with urea in the 60 - 70 C range which is considerably lower then what you are proposing. Iron is run at a higher temperature because of the higher
decomposition temp of the Fe(CO)6 and also due to the higher heat necessary for easy formation. Also I've seen the modification with molten urea
being pumped through a bed of hot nickel pellets.
Unionised, I think this process is not used because of the incredibly low yields, possibly long reaction times/wasted heat input (and it might not
even work).
SAM4CH, did you do any tests on your distillate to determine if it was indeed hydrazine? I am very interested in the quality/total quantity of your
product.
Not enough quantity!
SAM4CH - 18-7-2004 at 23:08
I wasn't lucky to get enough amount of hydrazine to make a test, I got few drops (about 1mL after 3hours) I am disturbed about result, So please
any one tell me about the best test for little drops.
I want to ask about the slowness of reaction and I think it might take 1day to make 10mL of hydrazine in this method. what about possibility of
explosion during distillation process with out inert gas and using glassware.unionised - 21-7-2004 at 13:19
This reference http://www.osha-slc.gov/dts/sltc/methods/organic/org020/org0...
should let you check for the presence of less than a microgram of hydrazine.
The colourimetric test isn't very selective and you can get misled into thinking there is hydrazine even when there is not.
The HPLC procedure is much more selective, but it's not a lot of use unless you have a laboratory.
It's sensitive enough that, when we used the same analysis at work, we were able to determine the concentation of hydrazine in commercial
ammonia (and that's not a lot of hydrazine).BromicAcid - 1-8-2004 at 17:04
Regarding the use of nickel powder in this reaction. I've found that nickel carbonyl is somewhat harder to make then I was initially lead to
believe. After finding a prep specifically for its production I have the following useful bits of information to add to my hydrazine excursion by
this method.
The slightest amount of oxygen negates the reaction and renders the nickel inactive.
Hydrogen sulfide stops this problem in even the smallest concentrations.
The reaction at 230C is only fast if the nickel is divided to almost pyrophoric levels, e.g, decomposition of formate or reduction of nickel
chloride with a sodium dispersion in kerosene.
A slight trace of mercury increases the reaction rate several fold, probably some in situ amalgam as usual.
Therefore I now have some modifications to my procedure.
Other interesting things I learned from the prep. Nickel carbonyl burns with a gray flame, and will explode without oxygen at about 350C according to
the following reaction:
Ni(CO)4 -----> Ni + 2CO2 + 2C
Leaving a dark cloud in its wake from finely divided nickel (which may be pyrophoric) and carbon. Nickel carbonyl readily decomposes above 50C and
the slightest contamination with air leads to:
2Ni(CO)4 + O2 ----> 2NiO + 8CO
The nickel floats around in the carbonyl as a 'flocculent' precipitate. Also carbonyls can have violent reactions with rubber stoppers.
