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Author: Subject: Formamide from urea
njl
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[*] posted on 27-3-2020 at 13:09
Formamide from urea


Background: Was bored. Have 5 pounds of urea. Have no formamide.

Theory: A well known route to ammonia gas/isocyanate is the reaction of NaOH with urea at elevated temperatures1. I thought maybe instead of the N-C bond being cleaved and the remaining (O=C-N) species being reduced to (O=C=N -), maybe the leaving amine could be replaced with a hydride (H-). A short period of research into this and related reactions returned no results whatsoever, which was discouraging. But, I have a lot of urea and time on my hands so I figured I'd give it a shot. The main products of the above reaction are ammonia and sodium isocyanate. I don't want sodium in my end product, so my next thought was to have the urea be in solution while a different redox reaction took place, the same way that the substrate in a Clemmensen reduction is not directly participating in the production of zinc chloride2.

First experiment: reduction with NaOH (aq)/Al (s)

Preface: This experiment wasn't well thought out. Looking back, I'm not sure what I expected the products to be or how I was going to know if the reaction worked as expected.

I took a few notes and wrote my thoughts and concerns down. First of all, the reaction between NaOH and aluminum is highly exothermic so maybe the reaction mixture would get hot enough to undergo the above reaction and just produce sodium isocyanate and ammonia1. Therefore, I reasoned that the mixture should be kept cool. Secondly, the only way I could tell that the reaction proceeded as I wanted it to was to smell the ammonia generated. However, since Ammonia is also a product of the isocyanate reaction, I would have no way of knowing if the reaction was in fact what I had hoped. Either way, I proceeded as follows:

Into a 100 ml beaker there was added 40 mL of purified tap water at
~20 C. To this, enough NaOH was added to form a saturated solution. The beaker was left to cool to rt, and then the aqueous NaOH was decanted into another 100 mL beaker while care was taken to leave the excess NaOH pellets in the bottom of the first beaker. To this solution, enough urea (pure but probably not dry) was again added to form a saturated solution. As before, this solution was decanted into a clean beaker to separate the aqueous reaction mixture from the excess urea.
An arbitrary amount of aluminum foil (maybe 4 grams) was torn into small pieces and crumpled into balls, that were then added to a test tube. Once half the foil was processed, I set aside the rest for later. Now, with gentle stirring, the aluminum in the test tube was gently added to the beaker and a glass stir rod was used to submerge each piece to ensure good contact. Nothing happened for a short period of time while the passive oxide layer was dissolved, and then the reaction continued as expected with the generation of small bubbles. Once a majority of that aluminum was reacting, I began preparing the rest of the foil in the same way. The remaining aluminum was added and the contents of the beaker allowed to cool. It should be noted that the reaction mixture never exceeded 30 C. Over the duration of the reaction, I routinely wafted air from above the beaker to try and see if any noticeable ammonia was being generated. At no point did I smell anything. Assuming this was a failure, I set the beaker aside and proceeded to the next trial.

Experiment 2: reduction with HCl (aq)/Al (s)

Preface: I had more hope for this trial as there was no sodium in solution and the products of the reaction would simply bubble out of solution as hydrogen gas/ammonia/water vapor.

To a 100 mL beaker there was again added 40 mL of ~31 percent HCl (w/w). Next, enough urea was added to the acid to make a saturated solution (much more than I expected despite urea's already high solubility in water, possibly due to the formation of even more soluble urea hydrochloride). The reaction mixture was then decanted into a fresh 100 mL beaker with the excess urea left in the first beaker. As before, approximately 4 grams of aluminum foil was torn into small pieces and crushed into balls, and stored in a test tube. About half of the foil was then gently added to the reaction mixture. The beaker was swirled and the floating foil balls were submerged with a glass stir rod. No observable reaction took place (I think the far more concentrated NaOH solution removed the oxide layer more quickly) for quite a while. However, once the reaction started (as indicated by the generation of bubbles on the surface of the foil) it became quite violent and splashed some of the contents of the beaker on to the wall of the beaker and my workbench. just as quickly as it took off, the mixture settled down although bubbles were still being generated. I took this opportunity to waft some air from above the beaker to try to smell if any ammonia was present. I did smell ammonia which was accompanied by the characteristic burning. In the interest of consistency, I decided to continue to add the foil, albeit more slowly. Over the course of maybe 15 minutes, foil was added in portions such that the reaction never entered thermal runaway and stayed relatively calm. I did not get the chance to smell any generated ammonia since I was cleaning up the mess from the first addition and drafting this writeup. After the last addition, the beaker was quite warm, and was therefore allowed to cool back to rt.

Discussion: I should say that I have no formal experience or access to analytical equipment. I just ordered my first lab stand after a few months of actively experimenting with home chemistry so maybe I'll be able to try this reaction on a larger scale and distill off the product (if there is any). I did this experiment on whim, and I didn't and still don't think it went as I had hoped. Formamide might be a minor side product or the urea in solution could be totally unaffected by the formation of the aluminum salts. Two important notes: before attempting the first experiment, I did a quick control with saturated NaOH solution and an appropriate amount of foil. The generated gas had no odor (as expected of hydrogen gas) but was extremely irritating. I have no rationale for why this is, as I have never heard of hydrogen gas being an irritant, but it is consistent with other times I've done that same reaction in class. Secondly and perhaps more importantly, my sense of smell is (to be blunt) kind of shit. Sometimes I can't smell anything, other times it's perfectly normal. It would be totally reasonable to argue that no ammonia was generated at all and that I was mistaken.


Notes:

1. Nurdrage video

2. It's not really correct to say that the substrate isn't participating in the reaction, but I just mean that the main reaction going on (the formation of zinc chloride through the oxidation of metallic zinc and reduction of HCl) is providing the means for the substrate to be reduced, so in that way the reduction of the substrate is secondary to the main reaction.
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clearly_not_atara
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[*] posted on 27-3-2020 at 15:50


The problem with basic reductions is that urea gets deprotonated and then it's hard to transfer electrons to a negatively-charged molecule. The problem with acidic reductions is that most strong reducing agents react with protons and uronium isn't that reactive.

In order to break this conundrum you might be able to complex urea with a Lewis acid and reduce the resulting complex. One such reducing system is Al/CrBr3 which generates AlBr3 and CrBr2 in situ. Unfortunately I know of no good way to make anhydrous CrBr3 other than the (highly exothermic) combination of the elements. The possibility of HCN formation in such a system should be considered.

EDIT: as with some other metal halides, it may be possible to make CrCl3 by rxn of Cr with CuCl2 assisted by ethyl acetate

[Edited on 27-3-2020 by clearly_not_atara]




[Edited on 04-20-1969 by clearly_not_atara]
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njl
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[*] posted on 27-3-2020 at 17:26


@Atara thank you for the feedback. I agree that reduction in basic conditions is unlikely to be a viable route. While it is also my understanding that a majority of strong reducing agents react with acids I don't see how that relates to the method I proposed. Could you elaborate? Also, do you have a ref for the urea complexing/reduction? Sounds interesting.
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clearly_not_atara
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[*] posted on 28-3-2020 at 08:16


The point is, in a system like Zn/HCl, there are two things that can be reduced:

2 H3O+ + Zn >> 2 H2O + H2 + Zn2+

or

2 N2H4COH+ + Zn + H3O+ >> H2O + NH2CHO + NH3 + Zn2+

But the first reaction is much faster than the second reaction. So you get mostly reduction of protons to hydrogen with little reduction of urea to formamide. Zn or Al or Mg doesn't matter much in this case.

I don't know anything about reducing Lewis acid-urea complexes honestly. It was more of a wild guess.

One article suggests that my guess was sort of on the right track, using LiBr as the Lewis acid and SmI2 as the reductant for tetrasubstituted ureas:

https://onlinelibrary.wiley.com/doi/pdf/10.1002/ejoc.2018007...




[Edited on 04-20-1969 by clearly_not_atara]
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Σldritch
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[*] posted on 28-3-2020 at 08:56


I did something similar a while ago. Not sure what i expected but i think i got Aluminium Cyanate. https://www.sciencemadness.org/whisper/viewthread.php?tid=15... May be useful to know for you.
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