Sciencemadness Discussion Board

Hydrazine

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Taoiseach - 11-2-2010 at 09:23

Damn I can't wait to know if this method really works :P

Just heat the damn stuff somehow... put a few crystals in a testtube and heat over a candle. Or use a spoon over a lighter... be creative ;)

Btw. the crystals on your last pics look much more like HS!



[Edited on 11-2-2010 by Taoiseach]

Myfanwy - 11-2-2010 at 10:04

im just puryfiying it. i dissolved the 27g in 250ml boiling water,filtrated and cooled down.
no crystalls appeared at 0°C.
then i boiled 125ml away and cooled down. there was even no crystalls.
just after adding a cord the first crystalls settle on the bottom.
did N2H6SO4 sublimates in hot water? that would be a shame^^but would be very strange for a sulfate.
i think in 6 hours the whole glass will be full of crystalls.
then i heat the hydrazine sulfate in a test tube.
-

the HS dont really want to recrystallize. need help!
last time the whole liquid became a white mass, after the impulse when i opened the glass.

now very less crytsalls have formed.
the soltuion is ice cold.


[Edited on 11-2-2010 by Myfanwy]

you can see on the pics.
less than 1g crystallized.
HS dont like me and wants to stay in solution, or it sublimates in the boiling water.
i shouldnt had dissolve it in boiling water :( :D

fail

[Edited on 11-2-2010 by Myfanwy]

max1.jpg - 38kBmax2.jpg - 30kB

entropy51 - 11-2-2010 at 10:16

According to Merck Index hydrazine sulfate is soluble in 33 parts of water, freely soluble in hot water. So I think that no more than 3.8 gms would stay in solution in the 125 mL of solution that you ended up with.

rrkss - 11-2-2010 at 10:31

Take a glass rod and scratch the bottom of your flask, maybe that will cause the crystals to crash out of solution usually works for me when I'm having problems triggering crystalization otherwise reheat the flask gently in a hot water bath and allow it to gradually cool preferably in the hot water bath without disturbance.

Myfanwy - 11-2-2010 at 10:42

good idea, but i got it now.
found that HS is nearly insoluble in Ethanol at 25°C, so i added about 300ml.

now i have it back!
If you having problems to crystallize the HS from the solution, add some ethanol and all will be fine =D

:P

[Edited on 11-2-2010 by Myfanwy]

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rrkss - 11-2-2010 at 10:51

Good idea using ethanol to crash it out of solution.

Myfanwy - 11-2-2010 at 13:35

i throw my HS away!
seems to be very unpure.:o

i heated this stuff, untill the glass from the test tube melts, but it doesnt melt..
BUT with sodium hydroxide a fuming liquid was prepared, that reacts violently with <25% HCl.

i have to try the whole synthesis again. i made something wrong.
can u guys try it , too?

i will make it again on tuesday.
im very dissapointed now.
sorry i failed it :D

but im sure hydrazine sulfate can prepared that way!



[Edited on 11-2-2010 by Myfanwy]

[Edited on 11-2-2010 by Myfanwy]

Attachment: hydrazine und HCl.wmv (335kB)
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woelen - 11-2-2010 at 13:55

I have to agree with the remarks of len2. I have done quite a few experiments with TCCA and in all situations, where acid is added, the free cyanuric acid is formed and this is only very sparingly soluble in cold water, but much better in warm water.

I'm very sure that indeed you made hydrazine or its sulfate, or even its isocyanurate, but I do not expect the product to be very pure. The reaction you show in the small video also occurs with other molten/hot liquids containing lots of NaOH. Just mixing conc. NaOH-solution and conc. HCl will give a very violent reaction as well (neutralisation of the acid and base is very exothermic).

Waht you could try is adding NaOH to the mix of urea and TCCA and trying to distill off dilute hydrazine hydrate. This distillation is somewhat risky though, because you'll have a very viscous mass which may be hard to boil evenly. Hydrazine hydrate is fairly stable, hydrazine is not. Always assure that water is present in sufficient quantities, otherwise you might have an explosion (anhydrous hydrazine is said to be very unstable, but I have no personal experience with the anhydrous compound, I only have hands-on experience with the hydrate and salts of hydrazine).

So I think you can make hydrazine with TCCA and urea, but isolation of this or the sulfate of it is a hard job. The probem is that hydrazine sulfate and cyanuric acid have similar solubility properties in water. Also on addition of a strong base they behave similarly (both dissolve in strong base).

Nicodem - 11-2-2010 at 14:44

Quote: Originally posted by woelen  
I have to agree with the remarks of len2. I have done quite a few experiments with TCCA and in all situations, where acid is added, the free cyanuric acid is formed and this is only very sparingly soluble in cold water, but much better in warm water.

I had many occasions to use TCCA at my job and I can assure you that it is quite stable in strong acids as long as there is nothing present that can reduce it or otherwise react with it. This is one of its features along its increased electrophilicity that differs widely it from others typical haloimides (for example, NCS is known to slowly decompose in strongly acidic media trough autooxidation). Like I said, I used TCCA/H2SO4 solutions several times and as long as there is no organic substrate present it will not get reduced to cyanuric acid (particularly if there is no water present to hydrolyse it to mixtures of chlorine oxides). I also used it for electrophilic chlorinations in dichloromethane with methanesulfonic or triflic acids as catalyst and no unwanted decomposition happens (besides the intended reaction). Maybe you should tell which acid you used, in what solvent, and what else was there present.
Quote:
I'm very sure that indeed you made hydrazine or its sulfate, or even its isocyanurate, but I do not expect the product to be very pure.

I'm a surprised at all the optimistic replies even though no evidence is presented. My first thought would be that he isolated Na2SO4 decahydrate. From the kewlish way Myfanwy describes his experiments it is impossible to reach any conclusion, except that his chances of dying from cancer some time in future are now enhanced, but nothing really about the chemical aspects. Unless he buys a termometer for up to 300°C (for a dozen well spent euros) and measures the melting point, or does any other qualitative test, there is not really much to say.

Judging from the smell? Take ammonia smell from urea hydrolysis and wishful thinking and there you have it, the perfect hydrazine smell! I'm not claiming that you can not make hydrazine from urea and TCCA, actually I think it is possible, but taking a scientific approach is more likely to give positive results. Myfanwy, it is nice that you are so motivated and eager to do experiments, but don't forget to keep it scientific so that others can benefit from it as well.

franklyn - 11-2-2010 at 15:35

@ Myfanwy

Even if unsuccesful we celebrate your effort. Nothing ventured , nothing gained.
If you know exactly what the outcome of what you are doing is going to be , then
you are not doing science , which by definition is investigation of what is unknown.
Edison performed 17.000 experiments to develop his incendescent lightbulb, his
token remark of which was invention is 1 % inspiration and 99 % perspiration.
Adversity and setback is most of experience.

.

woelen - 11-2-2010 at 23:53

Quote:
I had many occasions to use TCCA at my job and I can assure you that it is quite stable in strong acids as long as there is nothing present that can reduce it or otherwise react with it. This is one of its features along its increased electrophilicity that differs widely it from others typical haloimides (for example, NCS is known to slowly decompose in strongly acidic media trough autooxidation). Like I said, I used TCCA/H2SO4 solutions several times and as long as there is no organic substrate present it will not get reduced to cyanuric acid (particularly if there is no water present to hydrolyse it to mixtures of chlorine oxides). I also used it for electrophilic chlorinations in dichloromethane with methanesulfonic or triflic acids as catalyst and no unwanted decomposition happens (besides the intended reaction). Maybe you should tell which acid you used, in what solvent, and what else was there present.
Under the conditions you mention, indeed TCCA may be quite stable. I probably was not very clear in my formulation. I wanted to say that in reactions in which TCCA is used as oxidizer in aqueous medium in the presence of acid that cyanuric acid separates as an insoluble solid. I have not done experiments in non-aqueous solvents and then things might be completely different. I have read about TCCA being dissolved in acetone, but to my opinion that kind of experiments is too dangerous.
Quote:
I'm a surprised at all the optimistic replies even though no evidence is presented. My first thought would be that he isolated Na2SO4 decahydrate. From the kewlish way Myfanwy describes his experiments it is impossible to reach any conclusion, except that his chances of dying from cancer some time in future are now enhanced, but nothing really about the chemical aspects. Unless he buys a termometer for up to 300°C (for a dozen well spent euros) and measures the melting point, or does any other qualitative test, there is not really much to say.

Judging from the smell? Take ammonia smell from urea hydrolysis and wishful thinking and there you have it, the perfect hydrazine smell! I'm not claiming that you can not make hydrazine from urea and TCCA, actually I think it is possible, but taking a scientific approach is more likely to give positive results. Myfanwy, it is nice that you are so motivated and eager to do experiments, but don't forget to keep it scientific so that others can benefit from it as well.
Aren't you a little bit too harsh? I agree with you that things might look somewhat unscientific, but didn't we all start this way? I actually think that these attempts are nice, even if no practical results are obtained. He shows interest in chemistry, many other starters just go for the big BOOM and the F1r3 and Sm0ke. The reason why I think some hydrazine was made is because Myfanwy reported the formation of a silver mirror from AgNO3 and his white solid.

The only real concern I have is that Myfanwy should indeed be more careful with being exposed to the chemicals. If he is working like this with all his experiments and is exposed frequently, to all kinds of nasty fumes and gases, then he may experience bad health effects in the long run. So, my advice to Myfanwy is to scale down experiments to test tube size first and when things look promising then scale up to larger batches. Besides that, using smell for detecting stuff like hydrazine is not the wisest thing to do, it is better to find some chemical reactions which can detect its presence or even can determine the amount in a quantitative way.



len2 - 12-2-2010 at 02:06

Quote:


Also there is no need for acidic media



Right, and where do you think the proton to protonate the chlorinated acid into its reduced form is going to come from, given urea is oxidized to hydrazine?

Nicodem - 12-2-2010 at 05:17

I'm not sure I understand what you mean, especially the "protonate the chlorinated acid" part. Are you asking where the protons in cyanuric acid come from? If so, I already mentioned it above, from the reaction with urea which gives monochlorourea, in analogy with the known reaction of TCCA (and haloimides in general) with amides and sulfamides:

(CO-NCl)3 + 3 H2NCONH2 <==> (CO-NH)3 + 3 H2NCONHCl

The rearrangement however needs the presence of a base according to the classic (likely) mechanistic approach, which is furthermore supported by the relative stability of monochlorourea in acidic media.
For example, if the rearrangement happens according to the "aza-Favorsky rearrangement", you can check the depicted mechanism at the Wikipedia entry (just change the methylene group neighbouring to the carbonyl for the -NH2 and the -CHCl- group for the -NHCl). The formed N-aminocarbamate anion is unstable in neutral or acidic media and it decarboxylates to hydrazine.

len2 - 12-2-2010 at 19:45

Right, and where do you think the proton to protonate the chlorinated acid into its reduced form is going to come from, given urea is oxidized to hydrazine?

Nicodem - 14-2-2010 at 15:03

Len, there was no need to arrogantly copy-paste your question without answering my request for less vagueness. Particularly when I already answered it in all details (bold text part included!).
Just what proton you can not account for? If you write down both reactions that I discussed above (and their sum) you can see that there are no protons unaccounted for:

N-Chlorination step: (CO-NCl)<sub>3</sub> + 3 H<sub>2</sub>NCONH<sub>2</sub> <==> (CO-NH)<sub>3</sub> + 3 H<sub>2</sub>NCONHCl (12 hydrogens on both sides)

Rearrangement step: H<sub>2</sub>NCONHCl + OH<sup>-</sup> => H<sub>2</sub>N-NH<sub>2</sub> + Cl<sup>-</sup> + CO<sub>2</sub> (4 hydrogens on both sides)
-------------------------------------------------------------------------------------------------------------------------
Overall oxidation: (CO-NCl)<sub>3</sub> + 3 H<sub>2</sub>NCONH<sub>2</sub> + 3 OH<sup>-</sup> => 3 H<sub>2</sub>N-NH<sub>2</sub> + (CO-NH)<sub>3</sub> + 3 Cl<sup>-</sup> + 3 CO<sub>2</sub> (15 hydrogens on both sides)

I counted and recounted and I still can not account for your attitude.
Quote: Originally posted by woelen  
Under the conditions you mention, indeed TCCA may be quite stable. I probably was not very clear in my formulation. I wanted to say that in reactions in which TCCA is used as oxidizer in aqueous medium in the presence of acid that cyanuric acid separates as an insoluble solid. I have not done experiments in non-aqueous solvents and then things might be completely different. I have read about TCCA being dissolved in acetone, but to my opinion that kind of experiments is too dangerous.

But cyanuric acid also separates in organic solvents as well, but only after enough TCCA gets consumed in the reaction, so it is not like it is much different in this respect. This is actually a nice indication of the reaction progress, but only in solvents where TCCA dissolves completely. In acetonitrile, alcohols and acetone it is well soluble, but in CH2Cl2 it has lower solubility. The only difference is that in water cyanuric acid precipitates in larger crystals as dihydrate, while in organic solvents it drops out as fine white powder of anhydrous form. Only in reactions carried in H2SO4 the cyanuric acid side product does not precipitate until you quench the reaction by pouring it over ice - cyanuric acid is soluble in concentrated H2SO4! And of course in aqueous basic media where it is also soluble.
TCCA in acetone is no problem as long as the media is neutral (no acid or base must be present). You can make an experiment and put a little solution of TCCA in acetone in some dark place at room temperature and check how much time it takes for a precipitate to form. Obviously when using acetone as solvent one expects the reaction of TCCA to be much faster with the substrate in comparison to the solvent.
Quote:
Aren't you a little bit too harsh? I agree with you that things might look somewhat unscientific, but didn't we all start this way? I actually think that these attempts are nice, even if no practical results are obtained. He shows interest in chemistry, many other starters just go for the big BOOM and the F1r3 and Sm0ke. The reason why I think some hydrazine was made is because Myfanwy reported the formation of a silver mirror from AgNO3 and his white solid.

Yes, maybe just a little too harsh, but I was taking into account also his other posts in other threads. Kewlish or not, I don't want him to die poisoned by HCN, H2S, hydrazine or whatever just because he is unable to mature faster. OK, maybe for a teenager it is not easy to get motivated for science in the absence of such stuff as cyanides and other supermegacool dangerous stuff. I can understand that - I used to be young as well. But still he should realize that this is an amateur science forum where not all are pleased to read about such adventures and that some members expect a certain level is maintained. It is all about equilibrium - a little push here, a little push there, and he will get into the desired direction and maybe become a successful researcher some day in the future (if he survives long enough).
I also suspect there might have been some hydrazine formed there, but obviously it must have been very little or else he would have had less troubles crystallizing it out before all the Na2SO4.10H2O starts to (this salt is not particularly soluble in cold water). But since he does not report scientifically we don't even know if he added enough H2SO4 to protonate the eventual hydrazine. He used excess urea and a huge excess of NaOH. I think it would make more sense starting with stoichiometric ratios to avoid several practical problems and also because there is no need for using any excesses (I also do not see much point in adding gelatin in the first trials). He also was adding NaOH into the putative monochlorourea solution instead of the opposite (hydrazine can get oxidized by N-chloroamides, especially in basic media!).

len2 - 15-2-2010 at 17:51

Quote:


N-Chlorination step: (CO-NCl)3 + 3 H2NCONH2 <==> (CO-NH)3 + 3 H2NCONHCl (12 hydrogens on both sides)

Rearrangement step: H2NCONHCl + OH- => H2N-NH2 + Cl- + CO2 (4 hydrogens on both sides)
-------------------------------------------------------------------------------------------------------------------------
Overall oxidation: (CO-NCl)3 + 3 H2NCONH2 + 3 OH- => 3 H2N-NH2 + (CO-NH)3 + 3 Cl- + 3 CO2 (15 hydrogens on both sides)



So now you have protonation of your chloramine to give you the acid in basic aqueous solution.

Nicodem - 16-2-2010 at 14:33

Len, again you do your best to be incomprehensible. There are no chloramines involved. Do you perhaps mean TCCA or N-chlorourea? None of these needs to get protonated. And which acid do you mean? Cyanuric acid is nearly insoluble in water so it does not go the next steps. So you probably mean CO2? Well, the reason for that second reaction being written as it is, was in that I was not sure if the decarboxylation was a discrete reaction step or not, so I wrote the reaction as a neutral media outcome (which was just fine for I was answering your question about protons). Now I know better that the N-aminocarbamate salt formed during the rearrangement does not decarboxylate under the basic reaction conditions, but only after adding H2SO4 (or whatever acid of choice). So I was wrong and the second reaction should actually be two discrete reactions:

Rearrangement step: H<sub>2</sub>NCONHCl + 2 OH<sup>-</sup> => H<sub>2</sub>N-NHCOO<sup>-</sup> + Cl<sup>-</sup> + H<sub>2</sub>O
Decarboxylation step: H<sub>2</sub>N-NHCOO<sup>-</sup> + H<sup>+</sup> <=> H<sub>2</sub>N-NH<sub>2</sub> + CO<sub>2</sub>

I did a literature search on the topic of the rearrangement of N-haloureas into hydrazines and to my surprise realized that what appeared as a rather obscure reaction is not that obscure after all. In one instance it was even given a name, the Schestakow rearrangement (or Shestakov; after the author of the seminal papers). The reaction was found to proceed in three steps, the same ones I mentioned above, that is N-halogenation/rearrangement/decarboxylation (see J. Am. Chem. Soc., 76, 2572–2574). The N-halogenation reagents most commonly used are NaClO and NaBrO, but other halogenating reagents can be used as well. The rearrangement of N-chlorourea only occurred when heating under highly basic conditions (so I was wrong in saying that the reaction might work without using excess NaOH and alkali carbonates are not necessarily basic enough). Heat is not always used though. The decarboxylation of the so formed N-aminocarbamate salt requires neutralization, so no hydrazine forms until the final addition of the acid. N,N-Disubstituted ureas also undergo the reaction, so the mechanism appears to be that of a Hofmann rearrangement and not Favorsky (though this does not prove or disprove that both mechanisms are operating at the same time where possible, which is however limited only to N-monosubstituted ureas). I could find no examples of rearrangement of N,N’-disubstituted ureas which cannot undergo a Hofmann rearrangement, but could potentially undergo a Favorsky rearrangement. So either nobody bothered trying the reaction on N,N’-disubstituted ureas or the rearrangement cannot proceed through a Favorsky-like mechanism. This raises an interesting question and maybe when I’ll have some free lab time I’ll check the reaction on N,N’-dicyclohexylurea (I have some leftovers from couplings with DCC). In fact, a lot of work on using this reaction on substituted ureas was done due to the potential industrial use in the synthesis of carbidopa, so chances are that N,N’-dicyclohexylureas were simply neglected and thus never tried.

I currently do not have the time to acquire and read all the papers (some I do not even have access to), but for those interested, here are the references (I screened Beilstein, CA and Spresi).

Seminal papers and patent by Schestakow (the papers are in Russian):
DE164755
Zhurnal Russkago Fiziko-Khimicheskago Obshchestva, 35, 1903; 858.
Zhurnal Russkago Fiziko-Khimicheskago Obshchestva; 37, 1905, 5.
Other references for plain urea to hydrazine:
DE729105
J. Am. Chem. Soc., 76, 2572–2574.
Chemistry and Industry (London, United Kingdom),1954, 1452.
Journal of the South African Chemical Institute, 9 1956, 37.
Substituted ureas to substituted hydrazines:
Heterocycles, 12, 1979, 1571-1574.
Tetrahedron, 43, 1987, 891-894.
Indian J. Chem., A19, 1980, 825-828.
Chem. and Pharm. Bull., 31, 1983, 423-428.
Org. Prep. and Proced. Int., 17, 1985, 1-9.
Acta Chem. Scand., B29, 1975, 93-98.
Org. Synth., (1987) 173-182.
J. Chem. Soc. Perkin Trans. 1, 1990, 1319-1329.
Journal fuer Praktische Chemie (Leipzig), 316, 1974, 347-348.
Liebigs Ann. Chem., 1990, 949-952.
Journal of Organic Chemistry, 36, 1971, 1949-1951.
Patents: US5300688, ES2026759, HU9902484, PL162497, PL163125

len1 - 16-2-2010 at 17:56

Now I can agree with some of what you say. The reaction is indeed an analogue of Hoffman rearrangement (I have never seen it referred to by any other name). This typically occurs with hypochlorites although some chloramines work as well (reagents of type (RR')-NCl are called chloramines of which chlorimides to which TCCA belongs are a subgroup). (TCCA is not a good chloramine for this reaction for reasons below).

The protonation remark is not see if you can balance hydrogens, but a thermodynamic one. What drives the reaction of TCCA with water (or with amines in non aqueous media) is the formation of cyanuric acid, which can not happen in aqueous basic conditions. you rightly step away from that now, by dividing the reaction into two parts, the first part being in non-basic conditions. However this is all quite unnecessary, bacause acidic conditions, which is what I wrote about originally, speed up the reaction substantially - and in these the Hoffman rearrangemnt of urea doesnt work.

In basic media in which the Hoffman rearrangement of urea is normally carried out cyanuric acid and TCCA decomposes violently. Just add TCCA to NaOH solution and wait. This accounts for much of what this chap observed.

It is not true to say that hydrazine does not form until the neutralization stage - I can vouch for the fact that hydrazine is present as soon as the hypochlorite has reacted with the urea. This maybe is due to other reaction pathways.


For all these reasons direct reaction with TCCA is not a good way to hydrazine.

[Edited on 17-2-2010 by len1]

[Edited on 17-2-2010 by len1]

woelen - 18-2-2010 at 13:01

I purchased some hyhdrazine hydrate, it is from Acros Organics. The bottle is labeled "64% hydrazine" and this closely corresponds to N2H4.H2O. The label also says that the bottle must be stored in a refrigerator between 2 C and 8 C. The last thing makes me a little bit worried. Is this really necessary?

I have read many MSDS's and they all say that hydrazine hydrate is stable (as opposed to anhydrous hydrazine, which indeed is mentioned as unstable). Some MSDS's say that it must be stored in a tightly closed container at ambient temperature. What should I believe? Why would Acros suggest storage at low temperature? I don't like to have chemicals around which are unstable at room temperature.

[Edited on 18-2-10 by woelen]

len1 - 18-2-2010 at 18:38

At work we have a whole pile of chemicals that the bottle says need to be stored in that magical 2C- 8C: chlorotrimethylsilane, iodomethane, methyl borate, butyllythium, to name a few.

They are all air and moisture sensitive compounds, sealed therefore air tight. For a few of these, CH3ClSi for instance I think the temperature indication is just to keep the pressure down inside the bottle. They will still be stable at 15-25C.

turd - 22-2-2010 at 10:51

I have a plastic bottle of ~20% N2H4 stored at room temperature. It produces a few bubbles and slight pressure if left standing for a year or so. It says "repackage after 5 years" on the label. Subjectively I'd say it's about as stable as my 35% H2O2 solution, which also produces a few bubbles in a similar time frame.

12AX7 - 22-2-2010 at 13:48

And mix them and you'll get "a few bubbles" in a few microseconds :D :D :D

woelen - 22-2-2010 at 23:33

@12AX7: Actually, I tried that :D
I mixed 1 ml of 10% H2O2 with 0.5 ml of 64% N2H4 (added drop by drop). Nothing happens when this is done. The liquids mix and that's all.
Next, I added one granule of cobaltous sulfate (appr. 1 cubic mm, one crystal). At once a cloud of steam is produced and no liquid is left in the test tube.

watson.fawkes - 23-2-2010 at 04:15

Quote: Originally posted by woelen  
Next, I added one granule of cobaltous sulfate (appr. 1 cubic mm, one crystal). At once a cloud of steam is produced and no liquid is left in the test tube.
Why did you choose to use cobalt? Cobalt oxide is used for catalytic oxidation of ammonia to NO2 in nitric acid plants; it's the only catalyst other than platinum that's been viable in industry. So it doesn't particularly surprise me, but it does intrigue me, since this is not a gas-phase interaction.

woelen - 23-2-2010 at 05:45

The reason I used cobalt is that this is a good catalyst at high pH for redoxreactions in which H2O2 is involved. In such environments the cobalt is oxidized to hydrous Co2O3 and this in turns catalyzes the reaction. I have noticed this behavior in many redox reactions with H2O2 and since then it is my favorite catalyst when I want to try things like this.

Hydrazine by Anhydrous Chloramine & Ammonia Reaction

franklyn - 23-2-2010 at 22:24

The following supposition must be regarded as prospective research into unknown
but plausible investigation rather than elaboration of tried and tested preparations.

Direct reaction between Chloramine and Ammonia at the moment they are formed
is possible in a deep eutectic ionic liquid made of 1 mol of Choline Chloride with
2 mols of Ammonium Sulfate as a participatory solvent. 1 mol of Calcium Hypochlorite
( HTH is 65 - 70 % available Chlorine the balance is NaCl ) , and 1 mol of Lime ( CaO )
is then added in , rapidly mixing with a blender for around 5 seconds and left to stand
inside an air tight container. After a brief period of effervescence the expected result
should yield hydrated Calcium Sulfate and Hydrazine Hydrochloride.


(NH4)2SO4 + { Ca(ClO)2 + NaCl } => { NaCl } + CaSO4 • 2H2O + 2 H2NCl

HOCH2CH2(CH3)3NCl ( Choline Chloride ) => no reaction

(NH4)2SO4 + CaO => CaSO4 • H2O + 2 NH3
_______________

2 NH3 + 2 H2NCl => 2 H2N.NH2 • HCl


The occurrence of possible side reactions is limited by the minimal duration
of the sequence of events. Temperature , pH and trace presence of catalytic
metals promoting decomposition and dissociation, remain the antagonists of
the desired path. The reaction begins in a liquid state and may remain this
way ( given that hydrogen bonding sites prevail throughout providing miscibility
with Choline Chloride ). Without a means to absorb the heat released by the
concentrated reactants ( perhaps with an admixed inert salt ) highly vigorous
even violent activity can result from the quantity of gaseous formation. Given
the reactive nature of the precursors let alone the product , the yield cannot
be quantitative.
In a related process , according to :
Syntheses of Some N-substituted Hydrazines by the Anhydrous Chloramine Process
http://www.sciencemadness.org/talk/files.php?pid=131996&...
- Quote from pdf page 7 _
" Temperature of reactor wall does not seem to have any effect on the
Chloramine yield in the range 25 to 95 ºC, although Sisler et al (1954)
observed considerable decomposition of Chloramine when the reactor
walls are kept below the ambient temperature in the range 10 to - 30 ºC "
( my comment - counterintuitively contrary to what one would expect given
that chloramine is increasingly less stable with increasing temperature )

Dissolving the product in Methanol will precipitate the Ca(SO4) • (x)H2O , NaCl.
Filtering , then adding Chloroform to the solution should acquire a precipitated
Hydrazine Hydrochloride salt.


Choline Chloride - CAS 67-48-1
Physical Form: Crystalline
Color: White
Odor: Odorless
Molecular Weight: 139.62
Boiling Point: Decomposes > 302 - 305 ºC
Melting Point: 247 ºC ,
Water solubility: 650 g/L (20 ºC)
Soluble in , Alcohol , Chloroform , and Acetone.
Insoluble in Ether and Benzene
Acidic organic salt. Solution pH 6.5 - 8 ,
reacts as acid to neutralize bases.
http://www.inchem.org/documents/sids/sids/67481.pdf

Hydrazine Hydrochloride - CAS 2644-70-4
Chemical Formula: H2N.NH2 • HCl
Physical Form: Crystalline
Color: White
Odor: Odorless
Molecular Weight: 68.51
Boiling Point: Decomposes 240 °C
Melting Point: 89-93 ºC
Solubility In Water: 370 g/L (20 ºC)
See page 120 - Hydrazine Chemistry.pdf , attached below

Hydrazine Dihydrochloride - CAS 5341-61-7
Chemical Formula: H2N.NH2 • 2HCl
Physical Form: Crystalline
Color: White
Odor: Odorless
Molecular Weight: 104.97
Boiling Point: Decomposes 240°C
Melting Point: 198°C
Solubility In Water: Soluble
Specific Gravity: 1.42

Hydrazine - CAS 302-01-2
Physical Form: colorless, oily liquid, fuming in air.
Color: Colorless
Odor: penetrating odor resembling that of ammonia
Molecular Weight: 32.05
Density: 1.0036 (25/4°C)
Boiling point: 113.5°C (at 760 mm Hg)
Melting Point: 1.4 - 1.5°C
Solubility: soluble in Water, Alcohol, and Isobutanol;
Insoluble in Ether and Chloroform
Vapor density: 1.04 (air = 1)
vapor pressure: 14.4 mm Hg at 25°C
Explosive limits: 4.7 - 100% by volume in air
Flash point: 38 - 52°C (open cup)
Fire Hazard
It is a flammable/combustible material and may be ignited by heat, sparks, or flames.
Vapor may travel to a source of ignition and flash back. Container may explode in
heat of fire. Vapor explosion and poison hazard indoors, outdoors, or in sewers.
Runoff to sewer may create fire or explosion hazard. Vapors form explosive mixtures
with air. May continue to burn in the absence of air. Can catch fire when in contact
with porous materials such as wood, asbestos, cloth, earth, and rusty metals.
Incompatible with oxidizers, hydrogen peroxide, nitric acid, metal oxides, and strong acids.
Health Hazard
Organs affected include central nervous system; respiratory system; skin and eyes.
Chronic exposure in humans may cause pneumonia, liver and kidney damage. Liver
damage may be more severe than kidney damage. It is carcinogenic.
http://www.atsdr.cdc.gov/tfacts100.pdf

.

Attachment: Hydrazine Chemistry.pdf (709kB)
This file has been downloaded 1176 times

franklyn - 27-3-2010 at 15:27

Another possible hypothetical reaction scheme

The classic deep eutectic ionic liquid made of 1 mol of Choline Chloride with 2 mols of
Urea as a participatory solvent. 1 mol of Calcium Hypochlorite ( HTH is 65 - 70 %
available Chlorine the balance is NaCl ) , and 1 mol of Calcium Hydroxide ( Ca(OH)2 )
is then added in , rapidly mixing with a blender for around 5 seconds and left to stand.
After a brief period of effervescence the expected result should yield Calcium Carbonate
Calcium Chloride monohydrate, CO2 and anhydrous Hydrazine.

Vigorous foaming should be expected with the additional formation of Carbon dioxide.

(NH2)2CO + { Ca(ClO)2 + NaCl } => { NaCl } + CaCO3 + 2 H2NCl

HOCH2CH2(CH3)3NCl ( Choline Chloride ) => no reaction

(NH2)2CO + Ca(OH) => CaCO3 + 2 NH3
_______________

2 NH3 + 2 H2NCl => 2 H2N.NH2 •HCl

2 H2N.NH2 •HCl + CaCO3 => CaCl2 •H2O + CO2 + 2 H2N.NH2


Dissolving the product in Methanol will precipitate the CaCl2 •H2O and CaCO3. Filtering ,
then adding Chloroform to the solution should partition an upper Hydrazine phase.




[Edited on 28-3-2010 by franklyn]

franklyn - 4-4-2010 at 04:33

Yet another reaction scheme

Direct reaction between Chloramine and Ammonia at the moment they are formed
is possible in a deep eutectic ionic liquid as a participatory solvent made of 1.5 mol
of Choline Chloride with 3 mols of CaCl2 •6NH3 ( produced beforehand by introducing
ammonia into a sealed container holding anhydrous Calcium Chloride ).
2 mols of Trichloro-isocyanuric acid is then added into this , rapidly mixing with a blender
for around 5 seconds and left to stand inside an air tight container. After a brief period
of effervescence the expected result should yield Calcium Isocyanurate and Ammonium
Chloride and Hydrazine Hydrochloride.

The overall reaction is _

3 CaCl2 •6NH3 + 2 ( ClNCO )3 => Ca3{( NCO )3}2 + 6 NH4Cl + 6 N2H4 •HCl

HOCH2CH2(CH3)3NCl ( Choline Chloride ) => no reaction


The steps are _

6 NH3 + 2 ( ClNCO )3 => 2 ( HNCO )3 + 6 H2NCl

6 NH3 + 6 H2NCl => 6 N2H4 •HCl

3 CaCl2 + 2 ( HNCO )3 => Ca3{( NCO )3}2 + 6 HCl

6 NH3 + 6 HCl => 6 NH4Cl

____________________________


I now believe that my expectations for the results as given in the preceding two posts and
the third proposed method above are overly optimistic given that I have not accounted for
competing reaction pathways. The Chloramine as produced does not occur as it does in the
gas phase reaction of elemental Chlorine with Ammonia since Chlorine in those procedures
are not diatomic , sparing at least this initial formation of Ammonium Chloride.

2 NH3 + Cl2 => H2NCl + NH4Cl

Condensation of Chloramine with Ammonia produces Hydrazine and preferentially Ammonium
Chloride as well. Hydrazine will displace Ammonia from Ammonium Nitrate , it does not however
displace Ammonia from Ammonium Chloride , and that will not react further , sequestering it's
ammonia. As a result yield is likely to be half of what I projected unless twice the molar ratio
of reagent ammonia is provided.

2 NH3 + H2NCl => N2H4 + NH4Cl

Chloramine itself reacts with Hydrazine and is the major antagonist to the desired reaction.

H2NNH2 + 2 H2NCl => N2 + 2 NH4Cl

All indications are there is no way to produce hydrazine neat without extensive side products.
Clearly reaction kinetics largely determine the preparatory value of a prospective procedure.

.

NH2Cl + NH3 Process.gif - 99kB

Magpie - 6-5-2010 at 18:27

This is a huge thread for making hydrazine, and there are others. Why is it that I've only seen the Raschig method mentioned once? And there was no discussion of its merits or of anyone trying it. It's a very straightforward procedure, taking up less than one page in Brauer (p. 468).

Brauer (Raschig) specifies 20% aqueous ammonia, 1N (7.5%) NaOCl ,and H2SO4 as reactants. 1% limewater is used to complex metal ions. The equation presented is:

2NH3 + NaOCl + H2SO4 = N2H6SO4 + NaCl + H2O

I did note that the amount of NaOCl specified seems to be off by a factor of 10, however. Perhaps this is a typo?

Other than being able to use urea instead of 20% ammonia, which is less OTC, I'm puzzled as to why this method doesn't seem to be in use by members of this forum. Can anyone tell me why this is?

Formatik - 6-5-2010 at 22:54

Quote: Originally posted by Magpie  
This is a huge thread for making hydrazine, and there are others. Why is it that I've only seen the Raschig method mentioned once? And there was no discussion of its merits or of anyone trying it. It's a very straightforward procedure, taking up less than one page in Brauer (p. 468).

Brauer (Raschig) specifies 20% aqueous ammonia, 1N (7.5%) NaOCl ,and H2SO4 as reactants. 1% limewater is used to complex metal ions. The equation presented is:

2NH3 + NaOCl + H2SO4 = N2H6SO4 + NaCl + H2O

I did note that the amount of NaOCl specified seems to be off by a factor of 10, however. Perhaps this is a typo?


No, it's not. A similar ratio is seen in the OrgSyn method.
Quote:
Other than being able to use urea instead of 20% ammonia, which is less OTC, I'm puzzled as to why this method doesn't seem to be in use by members of this forum. Can anyone tell me why this is?


The percent yield is quite a bit better than some procedures that use ammonia (like the OrgSyn procedure).

It also could have to do with the volume of reagent (using less urea than ammonia solution).

Magpie - 7-5-2010 at 12:08

Quote: Originally posted by Formatik  
Quote: Originally posted by Magpie  

I did note that the amount of NaOCl specified seems to be off by a factor of 10, however. Perhaps this is a typo?

No, it's not. A similar ratio is seen in the OrgSyn method.


Thanks Formatik. I didn't even suspect that OrgSyn would have a synthesis for hydrazine sulfate. I thought that it would be commonly available from the US chemical industry. I guess when OrgSyn was started the US was almost totally dependant on Germany for such chemicals.

entropy51 - 7-5-2010 at 15:10

Quote: Originally posted by Magpie  

Thanks Formatik. I didn't even suspect that OrgSyn would have a synthesis for hydrazine sulfate. I thought that it would be commonly available from the US chemical industry. I guess when OrgSyn was started the US was almost totally dependant on Germany for such chemicals.
That's right, Magpie. See this link for the history.

kharof - 14-5-2010 at 04:31

i found these
http://www.freepatentsonline.com/pdf_collections_server1/usp...
ammonia + acetone + hypochlorous acid --> acetazine + H2SO4 --> HS
how i can to prepare hypochlorous acid?????????????
(Cl2 + H2O ---> HOCl + HCl i tried this but how to rremove HCl)
(Ca(OCl)2 + 2H2O --> Ca(OH)2 + 2HOCl i tried this but no smell chlorine and i remmember these hypochlorous acid smell chlorine)

i tried this:
ammonia + acetone + calcium hypochlorite and after a reaction i filtered it and added
H2SO4 produced precipitate large I heated it but not vaporize (i thought calcium sulfate) but when i added it into solution from NaOH smelled hydrazine)
HCl no thing (I thought should to precipitate CaCl2)
HNO3 (no thing)
then tried this:
ammonia + acetone + sodium hypochlorite and after a reaction a solution become yellow and evolve chlorine
then i added H2SO4 nothing
then tried this:
ammonia + acetone + H2O2 60% (i added sodium phosphate and acetamide(i not sure that is acetamide because i prepared it (i prepared it from CH3COOH + NH3CONH3 then reflux)))
and solution become boil when I added H2O2
and I tried it:i was sure these will failing
ammonia + H2O2 + gelatin+ heated in sealed bottle and after 2 minute my bottle exploded (pressure)
i prepared HS from method NH2CONH2 + naocl but yield is small and Nh3 + naocl but yield is very bad
my result: I need help

Magpie - 26-5-2010 at 14:46

I thought I'd post the results of two recent syntheses of hydrazine sulfate (HS).

My first synthesis utilized the chlorourea process as first expressed by garage chemist. Then ordenblitz developed a procedure from this. I used the ordenblitz procedure verbatim, except I substituted an equivalent amount of 12.5wt% NaOCl for Ca(OCl)2. My urea was recrystallized from urea fertilizer, which is dirt cheap. As ordenblitz stated, this procedure progressed quickly. Yield was 5.6g, or 33% based on the amount of Cl2 generated. A melting point of 254-255C was obtained, which agrees with the Baker Chemical MSDS value of 254C.

My second synthesis used the Hofmann reaction of urea as found in the procedure of Mr Anonymous. I used a scaled down batch size of 18.7%. Although more time consuming, the yield here was 31g, or 53% based on the amount of 12.5% NaOCl used. A picture of the HCl neutralization, which follows the generation of the hydrazine, is shown below.

It would be interesting to know which procedure is considered safer, especially since the Mr A procedure avoids a stable chlorourea intermediate.



neutralization during HS prep.jpg - 91kB

[Edited on 27-5-2010 by Magpie]

[Edited on 27-5-2010 by Magpie]

[Edited on 27-5-2010 by Magpie]

Rosco Bodine - 27-5-2010 at 00:48

The Mr. A procedure, which is actually an adapted Bayer patent process IIRC, is safer. It's probably the most straightforward synthesis of hydrazine there is. You know you want to scale it up though, probably in a 12 liter flask you can use a whole gallon of pool hypochlorite in one batch.
The stuff doesn't keep well so you may as well use it all at once. Beware the bubble monster ! :D

[Edited on 27-5-2010 by Rosco Bodine]

Magpie - 27-5-2010 at 06:37

Quote: Originally posted by Rosco Bodine  
The Mr. A procedure, which is actually an adapted Bayer patent process IIRC, is safer. It's probably the most straightforward synthesis of hydrazine there is. You know you want to scale it up though, probably in a 12 liter flask you can use a whole gallon of pool hypochlorite in one batch.
The stuff doesn't keep well so you may as well use it all at once. Beware the bubble monster ! :D

[Edited on 27-5-2010 by Rosco Bodine]


Yes, the Mr A procedure is quite straightforward, and I want to thank Mr A for doing the research work to optimize it.

In regard to getting rid of nearly 1 gallon of swimming pool bleach I must say that you have a point! Maybe I should start a business of sanitizing bathrooms, locker rooms, etc until I can get rid of it. :D

Rosco Bodine - 27-5-2010 at 07:43

You are welcome. I would reiterate that the Mr. A improved method is a very carefully worked out and many times tested and optimized procedure, which is reliable and confirmed multiple times, but is also a fickle reaction sensitive to small variables. The optimized reaction conditions window is narrow but is quite an easy target to hit if the process is followed exactly as described. Following the described process faithfully should reliably reproduce the stated yields. Varying the described process details by even a small amount will be something like fumbling around, awkwardly trying to find a keyhole in the dark. For all the doubters ...yeah well ....live and learn.

Anders Hoveland - 27-6-2010 at 13:32

I hesitate to repeat this, since it is already posted, but I feel it probably got overlooked by many so I am putting it here:
Making N2H4:
A buret slowly drips NaOCl bleach into a beaker. The beaker has two layers, a top NH4OH with some (NH4)2SO4 dissolved in it, and a bottom hydrophobic layer resistant to being oxidized by bleach. As hydrazine is formed, it immediately precipitates out as hydrazine sulfate (3.4g/100mL 25C) which falls into the bottom layer, where it is protected from further oxidation by the bleach. This would bypass the main difficulty of hydrazine synthesis: that bleach attacks the hydrazine formed faster than it makes chloramine. This might allow a higher yield of hydrazine, without having to use starch or gelatin. Ammonium sulfate solubility is 70.6g/100mL at 0C. Na2SO4 is 4.76 at 0C, not much higher than the hydrazine so you will likely get some of this precipitating out at the bottom too.
You might also see my post: "High Yield Hydrazine Synthesis?" which presents a completely different way to make dimethyl or tetramethyl hydrazine, where the hydrazine derivitives are protected from oxidation by bleach as soon as they are formed, so no starch or excess ammonia is required.

Alexein - 28-6-2010 at 13:19

Quote: Originally posted by Anders Hoveland  
I hesitate to repeat this, since it is already posted, but I feel it probably got overlooked by many so I am putting it here:
Making N2H4:
A buret slowly drips NaOCl bleach into a beaker. The beaker has two layers, a top NH4OH with some (NH4)2SO4 dissolved in it, and a bottom hydrophobic layer resistant to being oxidized by bleach. As hydrazine is formed, it immediately precipitates out as hydrazine sulfate (3.4g/100mL 25C) which falls into the bottom layer, where it is protected from further oxidation by the bleach. This would bypass the main difficulty of hydrazine synthesis: that bleach attacks the hydrazine formed faster than it makes chloramine. This might allow a higher yield of hydrazine, without having to use starch or gelatin. Ammonium sulfate solubility is 70.6g/100mL at 0C. Na2SO4 is 4.76 at 0C, not much higher than the hydrazine so you will likely get some of this precipitating out at the bottom too.
You might also see my post: "High Yield Hydrazine Synthesis?" which presents a completely different way to make dimethyl or tetramethyl hydrazine, where the hydrazine derivitives are protected from oxidation by bleach as soon as they are formed, so no starch or excess ammonia is required.


I don't think it would work, but i've been wrong before. Please post yields and pictures of your results, as well as characterization or confirmation experiments. I think silver nitrate works well to test hydrazine.

Rosco Bodine - 29-6-2010 at 02:23

Good not to confuse speculation with knowledge. Theory is just that. It has its usefulness for indication of what may be interesting basis for experiments to confirm or disprove what is predicted by theory. But one can get lost in theoretical ramblings which escalate to building further theory predicated upon earlier unproved theory as if it were an inconsequential missed step to establish the premise for advancing further. Scientific method is nuts and bolts systematic work and not a fiction novel added chapter by unproven chapter. Methods tested and proven valid and theory based proposed experiments are not to be confused. The first is a given and the second is a maybe. Sometimes the experimental and theoretical prediction of 1 + 1 = 2 is correct and sometimes it isn't. The experiment done and results determined accurately is what tells you if the theory was on track or off.

benzylchloride1 - 1-7-2010 at 19:28
Quote:

Beware the bubble monster !


I have had several experiences with the bubble monster. I ran this reaction several times at 1/2 scale in a 3L 5-neck flask with a mechanical stirrer several times without problem. The last two times i ran this synthesis, I lost much of the reaction mixture due to foaming upon adding the urea solution. I consistently got 50% yields when the reaction did not foam though.

Atofina PCUK cycle

gregxy - 9-8-2010 at 10:03

Has anyone tried this method to make hydrazine?

It uses NH3, acetone and H2O2

http://en.wikipedia.org/wiki/Atofina%E2%80%93PCUK_cycle

franklyn - 12-8-2010 at 10:07

Industrial methods are optimized for economy of scale
and are never one pot procedures.


Atofina Peroxide.gif - 23kB

The WiZard is In - 12-8-2010 at 16:04

Quote: Originally posted by Magpie  
This is a huge thread for making hydrazine, and there are others. Why is it that I've only seen the Raschig method mentioned once? And there was no discussion of its merits or of anyone trying it. It's a very straightforward procedure, taking up less than one page in Brauer (p. 468).

[snip]


A bit dated perhaps —

Audrieth and Ogg
The Chemistry of Hydrazine
John Wiley & Sons 1951

Describes the Raschig synthesis in some detail.

The WiZard is In - 12-8-2010 at 16:28

Quote: Originally posted by franklyn  
Industrial methods are optimized for economy of scale
and are never one pot procedures.


This the PCKU method is cover by US Patent 3 869 541. 1975.



Chemical-Industry-2.jpg - 206kB

Rosco Bodine - 18-8-2010 at 02:31

Quote: Originally posted by len1  
Quote: Originally posted by benzylchloride1  
I ran the hydrazine sulfate synthesis described on Sciencemadness that uses 10% NaOCl solution and urea, last year. This synthesis worked well and I obtained 185g of hydrazine sulfate. I am about to run the synthesis on a 1/2 scale in a 3 liter flask with mechanical stirring. The foaming problem seems to be the only issue with this synthesis; the first batch I attempted overflowed the reaction flask. At least it was in a fume hood, so the clean up was fairly easy. This was with the full scale procedure in the 3L RBF. I had the plastic funnel attached to the flask as stated. The second attempt was conducted in a 5 gallon plastic bucket and transferred to the 3L RBF after the foaming stage via a syphon. I obtained a good yield of hydrazine sulfate after neutralization, 185g. I am trying to see if this reaction could be scaled down without severly decreasing the yield. I synthesised 2,4 dinitrophenyl hydrazine, and 3,5-dimethyl pyrazole using some of the hydrazine sulfate. I plan to synthesize some luminol using readily available chemicals, including the hydrazine sulfate produced via this reaction.


The trouble with that synthesis, is that despite a great deal of detail its rather imprecise in its key point. For instance the volume of bleach is measured with great precision 1892ml, but then the bleach strength is quoted at 10%. Just a 1% change to 11% bleach strength alters the water content by 180ml, which makes a mockery not only of measuring the volume to that precision, but also the subsequent attempts to dissolve the NaOH in bleach, and the urea, and gelatine in a barely adequate amount of water.

There is the possibility that what seems to be a round figure of 10% in reality is accurately measured. That is unlikely, as the author does not seem to have performed bleach titrations. For instance he states that 10% bleach solutions decompose within a few weeks, to a couple of months. I have done bleach titrations, and a solution 11.5% strong, is still 8.2% even after 3 months non air-conditioned storage in temperatures which have exceeded 40 degrees for long periods of time.


Actually a 1% variance would equate more closely with 18ml of H2O than would it equate with 180ml of H2O, since the hundredth part (1%) of 1892ml is not the tenth part as your misplaced decimal point would indicate. You wouldn't want to be one to make a mockery of precision so let me help you stay straight on your arithmetic.

I have several very accurate instruments which I use when conducting the experiments which I sometimes do. If I state that I used a particular amount of a standard commercial item labeled and designated as a particular strength, then I stated the measured quantity accurately in all probability, although I am human and I may have inadvertantly stated 1892ml instead of 1900ml. However, it is more probable that if I stated 1892ml, then it was measured in a graduated cylinder of at least reasonable quality like Kimax or Pyrex class A glassware which would be accurate to plus or minus 2ml, so it may have actually been 1890ml or 1894ml which was used.

Not doing a bleach titration in order to satisfy your desire for greater precision should not be qualified as making the reported synthesis troublesome. Anyone is welcome to refine the details and embellish the process any way they wish, including doing bleach titrations if they think that is necessary to validate or make the work complete.

When doing bleach titrations, try to keep your decimal point in the right place and that way you will know better within plus or minus a power of ten what your bleach titration has revealed. Then you can be better equipped to quibble about what a difference one per cent would make.

On page 8 of this thread you will find the same file posted concerning the commercial bleach product and its analysis,
and although I did not do a chemical analysis of the bleach, I most definitely did check its density to have a reasonable certainty of what I was talking about and stating as accurately as possible here.


Attachment: bleach.pdf (80kB)
This file has been downloaded 920 times

len1 - 18-8-2010 at 06:32

Another day, another smart alec argument on SM. Why dont some of you guys get more interested in science than trying to score senseless points over each other while contributing nothing scientifically.

It is clear you are just clamoring an argument since I did state a change of 1% to 11%. Not by 1%. Its an additive change, and % is the appropriate unit. The statement is perfectly clear and there is nothing wrong with it.

Scientifically its also valid, since bleach, as stated in one of the posts can change substantally in concentration over a matter of weaks. Commercial bleach is not quoted to 1% accuracy, nor should it be due to the time variability. Hence an additive change of 1% is much more reasonable to consider than a fractional 1%. So the point is also valid scientifically.

Rosco Bodine - 18-8-2010 at 10:52

There is no desire by me to be a smartass or argue. It simply raised my ire a bit over the indignity that you would think or propose by inference that I am a careless technician who would not show ordinary diligence about understanding or justifying my numbers. I am not an incompetent or careless technician. So generally if I state a number, then it is reasonably accurate and reliable. I do try to contribute scientifically and that pursuit is not a game for me. If what I report is inaccurate and does not withstand scrutiny then I will assuredly correct any error when I learn of it.

http://www.youtube.com/watch?v=97WlNx-vJGs&fmt=18 Spiritual Vitamins

http://www.youtube.com/watch?v=ZEbeMvaVs68&fmt=18 You're All I Need to Get By

You're all I need to get by.
Like sweet morning dew, I took one look at you,
And it was plain to see, you were my destiny.

With arms open wide, I threw away my pride
I'll sacrifice for you, Dedicate my life to you
I will go where you lead, Always there in time of need


And when I lose my will
You'll be there to push me up the hill
There's no, no looking back for us
We got love sure 'nough, that's enough
You're all, You're All I need to get by.
You're all I need to get by.


Like an eagle protects his nest, for you I'll do my best,
Stand by you like a tree, dare anybody to try and move me
Darling in you I found
Strength where I was torn down
Don't know what's in store but together we can open any door
Just to do what's good for you and inspire you a little higher
I know you can make a man out of a soul that didn't have a goal
Cause we, we got the right foundation and with love and determination

You're all, you're all I want to strive for and do a little more
You're all, all the joys under the sun wrapped up into one
You're all, You're all I need to get by.

[Edited on 18-8-2010 by Rosco Bodine]

franklyn - 18-8-2010 at 13:08

Attachment: Hydrazine Sulfate preparation from Practical Inorganic Chemistry - Vorobyova.pdf (101kB)
This file has been downloaded 1072 times
Excerpted from Practical Inorganic Chemistry by Vorobyova.O.I., Dunaeva.K.M., Ippolitova.E.A., Tamm.N.S.
Cited here - http://www.sciencemadness.org/talk/viewthread.php?tid=9824&a...
relates to the preparation cited earlier in this thread by YeOldeImpurities
http://www.sciencemadness.org/talk/viewthread.php?tid=1128&a...

.

len1 - 19-8-2010 at 04:54

No I did not mean it like that. Everyones human. I make errors like that. Doesnt mean anything characterwise. And whats the point of me posting if i dont get criticism. it just has to be objective chriticism then its fine. People dont understand here, undeserved flattery is as good as condemnation. This was posted a while back. I actually did the experiment to see how it compared to my previous method. Got an explosion because my hypochlorite was 12%. hence the post

Rosco Bodine - 19-8-2010 at 11:32

Quote: Originally posted by len1  
I actually did the experiment to see how it compared to my previous method. Got an explosion because my hypochlorite was 12%. hence the post


Hmmm.....that is interesting .....but mostly because it is so highly unlikely

http://www.youtube.com/watch?v=i7RbZNT5A2A&fmt=18 Psychobabble

http://www.youtube.com/watch?v=NNiie_zmSr8&fmt=18 Eye In The Sky

[Edited on 19-8-2010 by Rosco Bodine]

len1 - 19-8-2010 at 14:09

Ah yes. Well now you are doing exactly what you accused me of. Playing the man and not the ball.

Rosco Bodine - 19-8-2010 at 15:40

Hey I'm having a little fun with a situation where you just keep digging a deeper hole on some bad numbers, and uh oh .....you dug slam to china where there
was a reported "explosion" which now makes the bad numbers good ? Gimme a break and just knock it off with the B.S. How about a detailed description of the process gone awry, so we can all wrap our heads around the "12% anomaly"
which you are now "reporting?" Because it seems pretty far fetched to me and I can guarantee you I'm not alone.

len1 - 19-8-2010 at 17:18

Just when it looked like you might be a reasonable guy, proof again that Im dealing with some nasty wacko. You are not alone, there are a number of you like that here. A testament to who you are is served by the inconsistencies of your posts and the sum total of what you have achieved. I think Ill pass on trying to explain anything to you, let your ignorance be your reward.

Rosco Bodine - 19-8-2010 at 18:32

Please do cite chapter and verse about whatever "inconsistencies" you allege exist about my posts. I will gladly correct or retract whatever error needs that attention.

Why do you make unfounded criticisms of good work that is accurately reported and well done, and then venture to think it is a "debate" that is occurring? I criticized your post which was inaccurate and misleading.

Your response is what is ignorant, not the person or the work you criticize incorrectly.

len1 - 20-8-2010 at 06:12

Has there been perhaps two Rosco's? One on this page, one on the one before, and one Mr Anonymous?

Rosco Bodine - 20-8-2010 at 11:45

Just one Rosco. No inconsistency. Mr. Anonymous was an alias used for e-mailed postings which were made before I became a forum member. Rosco is also an alias. Neither alias is a testament to who I am, neither is who I am relevant to the validity of information I have offered. Where are you going with this, you can't find a problem with the information I have shared, so now you have a problem with my screen name? How does any of this relate to your misplaced decimal and other problems with your own work in justifying incorrect criticisms
of mine ?

teresap989 - 10-9-2010 at 19:21

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



__________________
watch free movies online

S.C. Wack - 24-10-2010 at 17:25

You want more Rachig, I have more Rachig. Worth reading? Well...

Preparation of Hydrazine by Raschig's Method.
By REGINALD ARTHUR JOYNER.
J. Chem. Soc. 123, 1114 (1923)

"Estimation" rather than preparation would have been more precise, since none is isolated. The results reinforce why it is done the way it is.

Of all the books to be popular, Vorobyova? What you've discovered is another example of why it's pointless to post...people don't read.

Attachment: jcs_123_1114_1923.pdf (449kB)
This file has been downloaded 2135 times

verode - 14-11-2010 at 13:40

Quote: Originally posted by chemoleo  
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!

The Hofmann degradation is a reaction between an amide and a mixture of bromine and sodium hydroxide solution. Heat is needed. You may use urea and you get hydrazine:P

verode - 14-11-2010 at 13:42

The Hofmann degradation is a reaction between an amide and a mixture of bromine and sodium hydroxide solution. Heat is needed. You may use urea and you get hydrazine

Ketazine Process to make Hydrazine

NurdRage - 12-12-2010 at 21:59

I've had success with the Ketazine process for making hydrazine:

http://www.youtube.com/watch?v=UB7vwIFCnR0

For those of you that don't want to watch a 10 minute video, here is the write-up:
------------------------------------------------------------
Fumehood required!

Start with 250mL of ammonia and add to it 100mL of methyl ethyl ketone. Stir and then slowly add 1/4 mole equivalent of sodium hypochlorite based bleach. if using 10% bleach then about 186g is needed. If using 6% household bleach about 310g is needed.

As the bleach is added the reaction will heat up and bubble vigorously, slow down the addition if it's bubbling too much. Chilling with an icebath is helpful to keep the reaction under control. After all the bleach is added, keep stirring until the mixture stops bubbling. Then stop stirring and allow it to stand for a few hours or so until it separates into two clear layers. if an icebath was used it will take overnight.

The top layer is methyl ethyl ketazine. Separate it using a seperatory funnel or by careful decantation.

In a separate container, add 20mL of concentrated sulfuric acid to 100mL of water and stir. Then add the hot solution directly to the ketazine and keep stirring. The hydrolysis will generate crystals of hydrazine sulfate. Let solution cool to room temperature on its own and then filter off the crystals of hydrazine sulfate.

Acetone can be used in place of MEK, but you'll need to distill it off (use 50mL acetone in the reaction and distill 150mL of liquid) or extract it out with toluene into the acid solution.
----------------------------------------------------

The yield is around 30-60% depending on how slow and careful you are. It's not as good or as easy as the hoffman degradation method already posted in this thread. But i like the ketazine approach since its pretty robust and very forgiving if you mess up the concentrations or use alternative equipment or glassware. Purity is excellent, titration (with iodine) shows it to be 95%+ pure.

Maybe the hoffman degradation yield/purity can be improved by using a ketazine extraction step instead of treating with hydrochloric acid and sulfuric acids. I have not been able to test this myself since i'll be honest; i could never get the hoffman degradation to work reliably for me.

[Edited on 13-12-2010 by NurdRage]

BromicAcid - 13-12-2010 at 15:41

Great video, and I've always been interested in the ketazine process, nice to see it at work. Any clue if you have significant chloroform contamination in your organic phase considering you're adding hypochlorite to a methyl ketone?

NurdRage - 13-12-2010 at 16:05

To be honest i never actually measured that. I don't think it's an issue since the chloroform doesn't crystallize out with the hydrazine sulfate.

The side reaction might be responsible for the lower yields, but i never figured out how to inhibit it.

If you're looking for pure methyl ethyl ketazine then i suppose it would be an issue. in that case it can be remade from the pure recrystallized hydrazine sulfate by simply mixing it with aqueous excess sodium hydroxide and MEK.

[Edited on 14-12-2010 by NurdRage]

xxx - 29-12-2010 at 19:34

in patent 3869541 ketazine
why he used salt catalyst?
if anyone know?

UnintentionalChaos - 2-1-2011 at 10:07

I recently ran a hybrid hydrazine synthesis, combining the ketazine process and the Hoffman. I did a scaled-down Hoffman with a little excess urea and fresh 12.5% NaOCl. At the end, in lieu of acid neutralization, I added butanone and stirred for a while, then seperated off the organic phase. this was mixed with aqueous sulfuric acid and stirred. Heating was then applied and the butanone distilled off as it's azeotrope with water. The resulting solution was cooled and the product filtered and washed to give 52.5% yield of putative hydrazine sulfate. After drying, 77% of the butanone was recovered.

My question concerns the slight excess of urea. Does urea react with hydrazine to give aminoguanidine? Even if the reaction is slower, the hydrazine seems to very readily react with ketones and I'm worried that I may have quite some impurity in my final product. I could do an iodine titration or take a m.p. but I'd like to hear if my concerns have any basis first.

[Edited on 1-2-11 by UnintentionalChaos]

hydrazine sulfate...

lucky123 - 9-1-2011 at 00:54

I was just curious what can be made with this hydrazine sulfate? Primaries and secondaries or is it explosive by itself?

Random - 11-1-2011 at 14:41

Can I precipitate hydrazine with something other than sulfuric acid? Maybe hydrazine can be precipitated using potassium metabisulfite?

woelen - 10-2-2011 at 10:43

No, don't use potassium metabisulfite. What can be used though is NaHSO4 (pH minus for swimming pools). Dissolve this in water and add to the hydrazine solution. Your product, however, will be somewhat contaminated with Na2SO4 and/or NaHSO4. For most practical purposes, this is not a real problem.

Random - 14-2-2011 at 07:25

I thought about it as potassium metabisulfite should decompose to so2 and khso3 which would be oxidized by oxygen to KHSO4. Or maybe something like Na2SO4 would also work?

NurdRage - 14-2-2011 at 15:42

are you using the ketazine process or the hoffman degradation method?

Random - 15-2-2011 at 10:33

Ketazine process, it seems easier. Just I need a way to isolate hydrazine without sulfuric acid, though I can make above compounds if that would work.

UnintentionalChaos - 15-2-2011 at 11:40

Quote: Originally posted by Random  
Ketazine process, it seems easier. Just I need a way to isolate hydrazine without sulfuric acid, though I can make above compounds if that would work.


So you have butanone azine isolated? To hydrolyze that, you need a strong acid; bisulfate will probably not cut it. HCl would probably be fine, but you'll want to precipitate hydrazine sulfate afterwards with a bisulfate salt. The hydrochloride is quite soluble, IIRC

Random - 15-2-2011 at 12:21

Well, I am going to make methyl ethyl ketazine, just I wanted to be sure that I will be able to isolate hydrazine from that so it wont be waste of chemicals.
Now, I would then do that:

Isolate Ketazine layer, add HCl to hydrolize that (hydrazine hydrochloride in solution).
Mix potassium metabisilfite (decomposes to KHSO3 and SO2 in solution). Leave it in air to oxidize to KHSO4.
Add KHSO4 and SO2 solution to hydrazine hydrochloride and methyl ethyl ketone solution.
Hydrazine Sulphate should precipitatate. Now I need to filter it and the reaction is finished.

Now, is this right and would that work? By the way, what is the solubility of hydrazine sulfite, is it soluble or insoluble? Because I don't want to make mess because of SO2. Or maybe hydrazine sulfite is insoluble too. I know MEK bisulfite adduct is almost insoluble too.

If someone wants to know source of MEK, you can get it from some nail polish removers mixed with isopropanol, it shouldn't make any trouble in the reaction though, maybe somemore bleach should be used due to haloform reaction competing with isopropanol.

UnintentionalChaos - 15-2-2011 at 12:46

Quote: Originally posted by Random  
Well, I am going to make methyl ethyl ketazine, just I wanted to be sure that I will be able to isolate hydrazine from that so it wont be waste of chemicals.
Now, I would then do that:

Isolate Ketazine layer, add HCl to hydrolize that (hydrazine hydrochloride in solution).
Mix potassium metabisilfite (decomposes to KHSO3 and SO2 in solution). Leave it in air to oxidize to KHSO4.
Add KHSO4 and SO2 solution to hydrazine hydrochloride and methyl ethyl ketone solution.
Hydrazine Sulphate should precipitatate. Now I need to filter it and the reaction is finished.

Now, is this right and would that work? By the way, what is the solubility of hydrazine sulfite, is it soluble or insoluble? Because I don't want to make mess because of SO2. Or maybe hydrazine sulfite is insoluble too. I know MEK bisulfite adduct is almost insoluble too.

If someone wants to know source of MEK, you can get it from some nail polish removers mixed with isopropanol, it shouldn't make any trouble in the reaction though, maybe somemore bleach should be used due to haloform reaction competing with isopropanol.


Metabisulfite hardly oxidizes with any rapidity. That won't work. Sodium bisulfate is extremely cheap as pool pH down. In addition, bisulfite forms an addition product with MEK. if hydrazine sulfite is insoluble, it will be heavily contaminated with MEK bisulfite adduct. Isopropanol will be oxidized to acetone by hypochlorite (thereby lowering the yield of hydrazine), plus then some acetone azine can form (which is water miscible and will waste even more of your product.

Formatik - 15-2-2011 at 13:16

I tried out neutralization with sodium bisulfate using theoretical amounts and keeping the Na2SO4 just soluble with minimal water. This method looked promising, especially since the reaction between dilute H2SO4 and moderate aq. N2H4 is extremely violent (a lot worse than water and conc. H2SO4). But the yield I got was very low. I think compared to H2SO4, it is still not worth it because of great yield loss.

Here's what was done: 10.0g 35% N2H4 was added to a clear solution of 27.6g 95% NaHSO4 in 70mL of warm H2O. This caused slight fumes and appeared to cause only a mild exotherm. The reaction was expectedly many-fold weaker and far milder than with dilute H2SO4. After agitation especially, a glittering white precipitate formed. The solution cooled to 25 C eventually on standing, and at that time was filtered. And then washed, the amounts were approximate with a beaker, around like 20mL H2O for the first wash, then about 10mL for the second wash. Then air dried a few days. Yield: 8.60g N2H4 sulfate (60.5% of theory). Purity not determined.

Compared to neutralizations with dilute H2SO4 I've done, which are more predictable, don't need purifying washes and give a higher near theoretical yield (above 90%) even without further isolation from filtrate.

Probable reasons for the bad yield: around 21% of the yield was lost in filtrate and water wash, based on hydrazine sulfate solubility. The filtrate had that peculiar unique smell of hydrazine. More bisulfate could have been used. The composition of the bisulfate varies (it also can be a monohydrate). The impure pool bisulfate could use an assay for acid content. I can't comment on NaHSO4 usefulness in synthesis neutralizations. I would watch out for Na2SO4 crystallizing.

entropy51 - 15-2-2011 at 15:30

Quote: Originally posted by Formatik  
The composition of the bisulfate varies (it also can be a monohydrate). The impure pool bisulfate could use an assay for acid content.
I think the pool bisulate is usually the monohydrate.

Random - 16-2-2011 at 05:59

It looks like I will then need to make some H2SO4 to hydrolize ketazine. CuSO4 electrolysis seems easiest for dillute H2SO4.

NurdRage - 16-2-2011 at 06:52

Would you be able to find sulfuric acid based drain cleaner? Even if the acid is horribly impure the hydrazine sulfate can easily be recrystallized for maximum purity.

Random - 16-2-2011 at 10:22

I looked on many places, but I haven't found it. I was thinking about few other ways of production of it though:

1. Conversion of FeSO4 that is cheaper than CuSO4 to CuSO4 by some way like electrolysis. And then electrolysis of the resulting solution to H2SO4.
2. Dillute nitric acid (it would be mixed with some chlorine but it would only help) made from NO2 (Cu+HCl+Nitrate salt) and then react it with SO2 from metabisulfite + HCl.
3. SO2 + H2O2 --> H2SO4
4. Something like lead chamber process but I don't have enough lead to make it, it read it can be done using dillute concentrations in a PET bottle too.
5. React Cl2(g) + SO2(g) in water.

Every seems hard, though electrolysis of copper (II) sulphate with graphite anode and copper cathode and then filtering that should yield like 12% H2SO4 or something like that, which could be then used for Hydrazine.

UnintentionalChaos - 28-2-2011 at 09:58

Here's a video for anyone interested. I used a hybrid of the Hofmann and nurdrage's procedures to avoid exposure to free hydrazine. At no time during the synthesis was I able to detect any hydrazine smell (though I had carefully smelled a small amount the first time I ran the synthesis so I would know what it smelled like). The butanone azine has a significantly higher boiling point than free hydrazine so I was less concerned about exposure to that. The azine does have a strong and (to me) pleasant smell though.

Part 1: http://www.youtube.com/watch?v=JCrDttuw5co
Part 2: http://www.youtube.com/watch?v=URIAB32Q_JQ

[Edited on 2-28-11 by UnintentionalChaos]

Random - 28-2-2011 at 10:43

You have awesome channel and I see you often post your videos, I think those are really nice so keep up the good work :D

That synthesis with urea looks a lot better than ammonia, because urea is cheaper and more available, also there is not excess water.
Now I have everything what I need for ketazine synthesis, later I could hydrolise that to hydrazine sulphate, do you know if MEK ketazine is very toxic and dangerous to store? I actually didn't know that hydrazine was so toxic, I am afraid of nitrogen oxides and such things, while I thought hydrazine wasn't dangerous. It looks I was wrong.

NurdRage - 28-2-2011 at 10:49

Great work!

hydrazine is now completely within reach of everybody. The easy, low equipment requirement, and high yields of the original hoffman process, the high efficiency and materials savings of the hybrid process, and the highly robust (although expensive/labor intensive) ketazine process.


@random

I'm not certain about long-safety and toxicity of the azine. but the azines themselves are somewhat safer to handle than hydrazine since they are less volatile and not explosive.

sigma aldrich sells azetone azine directly so you might want to check the MSDS.

[Edited on 28-2-2011 by NurdRage]

Magpie - 28-2-2011 at 11:05

Beautiful presentation! This is great as no fume hood is required. I like the presentation of the Hofman degradation mechanism- very well done!

Unusual procedure hydrazine

AndersHoveland - 2-3-2011 at 16:47

Here are three lesser known procedures for preparing hydrazine:

USING OXIME PROTECTION
250mL ammonium hydroxide solution (20-30%)
100mL methylethylketone (MEKP)
0.25 moles sodium hypochlorite
(186g of 10% conc, or 310g of 6% conc solution)
until bubbles stop, about 10 minutes
oily layer will gradually form on surface, use separatory funnel to remove top layer

separately prepare a solution of:
100mL water
10mL 20%conc sulfuric acid
then pour in the oily layer from first preparation. The
solution becomes cloudy. filter out hydrazine sulfate

The following optional adjustments will improve yields: add the bleach very slowly using an addition funnel, cool the reaction with ice bath, boil the MEKP off before separating out the hydrazine sulfate. Up to 60% yield (based on initial hypochlorite) could be obtained, this is considered extremely good

USING UREA INSTEAD OF AMMONIA
There should be between 0.8 to 1 mole of hypochlorite for each mole of urea. For the one mole ratio of NaOCl, the yields were 50% with respect to urea, and 50% with respect to NaOCl. For the 0.8 mole ratio, the yields were 48.5% with respect to urea, and 60% with respect to NaOCl. For best yields, the hypochlorite solution should contain at least 3g of sodium hydroxide for every 0.8 moles of sodium hypochlorite. Solutions of sodium hypochlorite and urea were first both cooled to 4degC, then mixed. A glue solution was added, 0.5g of glue for each mole of sodium hypochlorite. The mixture was heated to 90degC, then allowed to cool. This exact temperature is very important for good yields from the reaction. Sulfuric acid was added into the solution so that the concentration of acid in the solution reached 20%. Hydrazine sulfate precipitated and was filtered out, then washed with alcohol and dried at 100degC. Note that the reaction cannot be done in a steel pot, as this catalyzes the decomposition of hydrazine.

OXIDATION WITH PEROXIDE
Hydrogen peroxide can be used to oxidize ammonia, with methylethyl ketone being used to protect the hydrazine as it forms. Acetamide, with some Na2PO4H added, must be present to activate the hydrogen peroxide, because otherwise the hydrogen peroxide and ammonia will not react. E. G. E. Hawkins, J. Chem. Soc. C, 2663 (1969). First, under the fairly alkaline conditions, ammonia condenses with the ketone to form an imine. The acetamide forms intermediate iminoperacetic acid, which then oxidizes the temporary imine to form an intermediate oxaziridine (which has a triangular ring with a C, N, and O atom. This ring immediately oxidizes an NH3 to form a hydrazone, which then condenses with another methylethylketone, protecting the hydrazine from further oxidation. The layer of methylethyl hydrazone separates out as a separate upper layer, which is decanted off.

NurdRage - 2-3-2011 at 18:01

Ya know... your first method would have been cooler if you didn't blatantly RIP-OFF my video, i even specified the exact same quantities....

Although the amount of sulfuric acid you specify is wrong and wouldn't yield very much hydrazine sulfate.

[Edited on 3-3-2011 by NurdRage]

bbartlog - 2-3-2011 at 18:04

So, where did you get the first procedure? When I see MEKP and sodium hypochlorite mixed, I tend to assume a haloform reaction will take place. Does the excess of ammonia mitigate this somehow? Or does the hydrazine reaction just take place faster? In short, are we supposed to believe that those who go to the trouble of doing this in two steps (make hydrazine, *then* add MEK) are just making life difficult for themselves?

(edit): actually I see that you've just copied NurdRage's method from upthread (and from his video as he notes). Interesting. Sounds like the hydrazine reaction outruns the haloform reaction.


[Edited on 3-3-2011 by bbartlog]

UnintentionalChaos - 2-3-2011 at 18:13

This is AndersHoveland...

I believe he was banned last time for spewing inane procedures that he had pulled out of his ass (often that made no sense) all over the forum. It looks like he's on his way to a repeat of that fate.

NurdRage - 2-3-2011 at 18:18

Yeah....

His second method is so vague as to be useless. If he means to replace the ammonia with urea from the ketazine process then i suspect its a blatant fabrication.... because i actually TRIED very hard to make a urea-ketazine process where i add the MEK and the urea together first... lets say i got all sorts of fun organic products except hydrazine.

i guess though he ripped another procedure, not a MEK based one. Those numbers sound familiar, i just can't put my finger on where.



[Edited on 3-3-2011 by NurdRage]

madscientist - 2-3-2011 at 18:51

Quote:
I believe he was banned last time for spewing inane procedures that he had pulled out of his ass (often that made no sense) all over the forum. It looks like he's on his way to a repeat of that fate.


Yep.

AndersHoveland - how many times do we need to tell you to use references? You've already long established yourself as unreliable. Everything you post needs to be verified.

Rosco Bodine - 5-3-2011 at 19:14

It seems it's going to be a long wait for me to see any truly significant improvement posted which goes beyond the improved method of
Mr. Anonymous .....

http://www.youtube.com/watch?v=uPHmsEoLsL0&fmt=18

NurdRage - 5-3-2011 at 20:00

The hybrid process was just posted by unintentionalchaos, it greatly reduces acid wastage and is partially recyclable. I think that's a significant improvement.

Rosco Bodine - 5-3-2011 at 23:47

Acid wastage? Yeah okay if you say so. I see a lot of time wastage from added complication, a lot of unnecessary glass plumbing, gradual additions which are yield crippling for what is an inherently "mass reaction" optimized situation
where reactants for the rearrangement need to be added in a lump to react en masse. The yield is likely being reduced by several variations on the process. The yield probably suffers specifically from scaling downward because of an unfavorable thermal dynamic, because it is very likely the reaction benefits from a sort of "critical mass" for the reaction quantity which enables it being self-driven by its own exotherm. If the scale and method I described was followed quite precisely in all details, there should be no reproducibility issues for independent confirmation and verification. The variations on a process susceptible to variables is what is introducing variables which are causing the yield problems. Following precisely the procedure I described, not on a different scale or with any variations or modifications whatsoever, should eliminate yield reproducibility issues. I still absolutely unequivocally stand behind what I reported in the improved procedure
http://sciencemadness.org/talk/viewthread.php?tid=757#pid653...
with regards to all of the described details of the process
and the reported yields. For the dubious, I would say
carbon copy, explicitly duplicate the process I described and the results should be consistent with what I reported.
If you do anything differently scalewise or otherwise and
then don't get as good yield, then logically it is probable that what was different is the reason for diminished yields for whatever technical explanation there may be. Analysis of the variables is an entire project by itself. However for the
scale and for the process I described, that information is reliable. The half gallon of pool chlorinator batch size would probably work fine in a 6 liter Erlenmeyer or a 6 liter Florence
which should contain all the foam without need for an overflow funnel. A full gallon of pool chlorinator batch size
could probably be done in a 12,000ml flask without overflow, but I can't guarantee what would be the foam dynamic there because I haven't tried it. A full gallon batch size should go fine in a 22liter or three gallons in a 50liter reactor. Have a ball with it if you have some process competence. But a closed garage without specific process ventilation is probably not a proper environment for this kind of experiment.

[Edited on 6-3-2011 by Rosco Bodine]

NurdRage - 6-3-2011 at 08:10

I'm not sure i see what you're specifically arguing.

Are you implying his method does NOT reduce acid wastage?

It looks to me you're arguing something else other than what i am claiming (reduction in acid usage)... i'm not sure what exactly.

If you can tell me in one line exactly what it is, i can give a more appropriate response.

Nicodem - 6-3-2011 at 08:48

Quote: Originally posted by AndersHoveland  
OXIDATION WITH PEROXIDE
Hydrogen peroxide can be used to oxidize ammonia, with methylethyl ketone being used to protect the hydrazine as it forms. Acetamide, with some Na2PO4H added, must be present to activate the hydrogen peroxide, because otherwise the hydrogen peroxide and ammonia will not react. E. G. E. Hawkins, J. Chem. Soc. C, 2663 (1969). First, under the fairly alkaline conditions, ammonia condenses with the ketone to form an imine. The acetamide forms intermediate iminoperacetic acid, which then oxidizes the temporary imine to form an intermediate oxaziridine (which has a triangular ring with a C, N, and O atom. This ring immediately oxidizes an NH3 to form a hydrazone, which then condenses with another methylethylketone, protecting the hydrazine from further oxidation. The layer of methylethyl hydrazone separates out as a separate upper layer, which is decanted off.

I was so shocked that you actually included a reference in that delirious post that I actually verified it (DOI: 10.1039/J39690002663). As expected, it has nothing to do with hydrazine synthesis and oxaziridines are not claimed as intermediates. That paper is about oxidation products of cyclohexanone/ammonia with H2O2 and no hydrazines, hydrazones or hydrazides are mentioned among the products (the products are mainly those derived from some sort of a Bayer-Villiger-like reaction on the cyclohexanone derived imine).
I can guess though that aimed at perverting the otherwise known oxaziridine based syntheses of hydrazones: Synthesis, 1991, 327-341, DOI: 10.1055/s-1991-26459 (attached).

I would appreciate if you would limit yourself to only post truth and nothing else but the truth - WITH TRUE REFERENCES!

Attachment: Electrophilic Aminations with Oxaziridines.pdf (1MB)
This file has been downloaded 2931 times

[Edited on 6/3/2011 by Nicodem]

Rosco Bodine - 6-3-2011 at 13:08

Ever heard that old phrase "penny wise but pound foolish" ? Maybe there is something of a mystery solving nature there in that cogent old phrase about economy, useful for decryption of my mysterious comments.

UnintentionalChaos - 6-3-2011 at 13:53

Quote: Originally posted by Rosco Bodine  
Ever heard that old phrase "penny wise but pound foolish" ? Maybe there is something of a mystery solving nature there in that cogent old phrase about economy, useful for decryption of my mysterious comments.


You have noticed that this is an amateur science forum, not a manufacturing plant, right? Most members here might have use for at most, a few mol of hydrazine sulfate, and almost surely not all at the same time. Since this is a hobby, they are not paid for their time and generally time spent doing chemistry is enjoyable to them.

I do not own any glassware larger than a liter, nor do I often have any use for such. For those doing such smaller scale organic chemistry, the glassware I used is most likely at the ready. 3 and 4L flasks and beakers are not. I went out of my way to use highly pure chemicals for this prep to avoid any chance of metal ion contamination in the product. It takes an obnoxious amount of time for me to distill azeotropic HCl and remove the iron contamination or whatever it is that gives it that lovely yellow color and I am not keen on throwing away several mole equivalents for direct neutralization.

Additionally, the smell of the azine is rather nice. While not healthy, I'm sure, it's worth a whiff.

I think you're just bitchy that nobody has given your ego a handjob lately over the original prep, but that's just my interpretation of the whole thing.

Rosco Bodine - 6-3-2011 at 15:02

If everything you were doing was right on track then where's the beef, where is the yield improvement or the process enhancement or simplification? It isn't there, so what exactly is the improvement? I haven't asked or been paid for my time sharing some insight into processes that I do understand with others, some of whom take offense at the theory or implication they don't know it all already .....
but who obviously know enough to correct or criticize or me, and if I'm actually wrong then I can appreciate the information. I don't need my ego stroked. I would settle for the respect of a presumption on the part of others that I do have two brain cells to rub together and make a spark.
Chemistry has dynamics about reactions which are real. You can accept that or pretend it isn't so and that variables don't matter. I have read posts that question the veracity of what I have reported because of variations by others on a process that was accurately described one way, being done a different way by others then seeing inferior results ....
Well Duh ....go figure why. Sure there is more than one way to skin a cat. That doesn't make plan B,C,D ect. necessarily
an improvement just because they are different from plan A.

NurdRage - 6-3-2011 at 15:40

I still don't know what he (rosco bodine) is so pissed off about. Can someone else tell me what's offended him so much? I wanna fix it, but i don't even know what the issue is.

@UnintentionalChaos have any plans to try the ammonia based ketazine process? Even on a small scale. I'd like to see how reproducible it is for others.

Rosco Bodine - 6-3-2011 at 16:20

I am not pissed. HCl is a cheap reagent so the acid wastage stuff is bs.
Using a ketazine variant (benzaldazine) isolation method was posted early in the thread, a German patent to a Russian named Shestakov. So this is not real news
except for the use of MEK being possible.

The interesting thing about a ketazine isolation method for hydrazine is a different matter.

What would be interesting is to see what may occur upon nitrosation of the ketazine in methanolic sodium hydroxide with isopropyl nitrite, or by nitrous gas nitrosation perhaps in the presence of isopropyl alcohol forming the IPN in situ.
Glycerin, ethylene glycol, or butanol in small amount may work better than isopropyl alcohol due to the less volatile nitrite which would function as a regenerable nitrite reagent in the system. Perhaps this could shortcut the route to sodium azide. Now that would be an improvement for a process whose eventual target product is sodium azide.
The nitrous gas I am contemplating would be by the method of Griess and Microtek using starch to decompose warm moderately concentrated HNO3. I think I recall from somewhere that paraformaldehyde can be used in the same way as starch for decomposition of HNO3.

Go back and check page 5 halfway down the page
http://www.sciencemadness.org/talk/viewthread.php?tid=1128&a...

The patent is here
DE164755 Schestakow , Hydrazine by Urea Hypochlorite.pdf (157kB)
http://www.sciencemadness.org/talk/files.php?pid=43024&a...

Also some commentary by Nicodem
http://www.sciencemadness.org/talk/viewthread.php?tid=1128&a...

[Edited on 7-3-2011 by Rosco Bodine]
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