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Author: Subject: Optimizing Henry reaction conditions (substituted benzaldehydes + nitromethane)
happyfooddance
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[*] posted on 5-7-2018 at 12:26


Quote: Originally posted by alking  
PEA should not smell much due to the high bp.


This makes no sense. Phenethylamine salts fume profusely when heated to their melting point. Additionally, PEA HCl smells quite a bit at room temperature. Most importantly, the boiling point of PEA freebase is lower than the melting point of the the hydrochloride salt. So when you free the base with KOH in the melt, your PEA will be superheated and boiling. On top of it all, adding KOH is probably exothermic in this case.

So I would suggest do it in a hood, or outside, and try it on a small scale first (like any sensible chemist would).
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[*] posted on 6-7-2018 at 08:05


I thought you were asking if PEA Freebase smells at RT or with use in this reaction, not if you boil it off lol. Of course it would smell a lot if you're boiling it.

At RT it has a faint odor, not anything to worry about. You don't need to melt the HCL salt or superheat it. The reaction at RT/moderately heated would be exothermic, but not anything approaching the bp. I've done this with just enough water and heat to get it to react efficiently, separated the PEA, and then dried it with Na2SO4 to pull any water out. It worked fine and there's less smell than a can of tuna.
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[*] posted on 7-7-2018 at 18:33


To all involved:
I took into consideration that PEA HCl liquifies at at greater temp then the freebase will volatilize. I answered this issue with Melgars original info. Using all modern safety and regulations... And INSIDE AN APPROVED FUME HOOD, I made a 20molar KOH solution, to a reaction vessel( 150x25mmborosilicate) added a 1/20 molar qty of PEA HCl. I then added 1.5 Mol qty of 20m solution. It slowly and if anything endothermically reacted. It got cold and separated into a thick salt layer on bottom and lighter amine on top. The melting of an HCl with a higher Bp then adding a salt to form a lower bP amine sounded dangerous. I’m happy to report in essence using a saturated KOH solution works as well as liquid HCl salt
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[*] posted on 7-7-2018 at 19:21


I had attempted to utilize Melgars ratios for large single crystals. I will admit it was ambiguous as fuck add 50%IPrOH, 25%1- butanol, 25%GAA. No mention of nitroalkanes besides “below 20% excess”. I honestly hold all of the old members in extremely high regard but some things do not lend themselves to repeatability. I’m not trying to mock him at ALL but I think all scientists published or not need to ensure their methods are repeatable and easily understood hence moles and shit like that. While here in the US our research stations crumble around the fume hoods for lack of funding
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[*] posted on 7-7-2018 at 19:51


I find that the refluxing acetic acid + ammonium acetate method is more reliable than the alcohol / base method. Nitromethane is cheap enough that a 5x excess can be used, in which case yields are quite good (typically 75–85%).

So far I have not attempted reductions of any nitrostyrenes. However, the Zn/HCl method seems like the most useful one, given that LiAlH4 is much harder to get (and also pretty dangerous).

[Edited on 7-8-2018 by Metacelsus]




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Melgar
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[*] posted on 19-7-2018 at 08:45


For anyone frustrated by all the missing details in my writeup, I guess the reason for that is that it wasn't really supposed to be a writeup. It was more like "I was experimenting with synthesizing nitrostyrenes using various solvent systems, and seem to have figured out something new, though I'm not sure what. Anyone feel like having a discussion about this?"

Quote: Originally posted by happyfooddance  
Quote: Originally posted by alking  
PEA should not smell much due to the high bp.


This makes no sense. Phenethylamine salts fume profusely when heated to their melting point. Additionally, PEA HCl smells quite a bit at room temperature. Most importantly, the boiling point of PEA freebase is lower than the melting point of the the hydrochloride salt. So when you free the base with KOH in the melt, your PEA will be superheated and boiling. On top of it all, adding KOH is probably exothermic in this case.

So I would suggest do it in a hood, or outside, and try it on a small scale first (like any sensible chemist would).


Okay, so you know how I said that I melted the PEA base and added KOH? Not only does this produce KCl and PEA base, it also produces substantial amounts of water. Since water has a much lower boiling point than PEA base, you will indeed get substantial amounts of water evaporating, and this keeps everything well below PEA base's boiling point. Not only is there water forming from the reaction, but KOH is always going to have a significant fraction of its mass be water initially, so this isn't something you really have to worry about.

I prefer doing it this way because the salt layer ends up as a slushy consistency that's more solid than liquid. You can just pour the PEA base off straight from a beaker, and the K-salt layer will stay stuck to the sides. This also combines several steps, and eliminates the need for titrating your KOH to figure out how much water it's absorbed and how much of it has converted to the carbonate salt. This means I can use my "old" KOH, in which the flakes have a powdery white carbonate coating from spending too much time exposed to air. It's always nice when you can find a use for your less pure reagents.

Anyway, you just add a moderate excess of KOH beyond what would be needed to neutralize the HCl salt. KOH is a strong desiccant too, so whatever doesn't react with the amine HCl salt will pull water out of it, which is important because PEA base is very soluble in water. I should also mention that I heated the vessel that I did this in, to just over water's boiling point, which sped things up a lot. Then I let it sit for a few hours to cool off, and give the K-salt layer sufficient time to desiccate the PEA base.

PEA HCl does not smell at all, or at least mine doesn't. This would be expected, since ammonium chloride doesn't smell, and neither does any pure amine HCl salt I know of. Any smell is almost certainly due to impurities. The free base does smell, although it's not particularly strong or unpleasant. Lots of things smell far below their boiling points or melting points. Vanillin, for example.

Quote: Originally posted by Gl3n  
I had attempted to utilize Melgars ratios for large single crystals. I will admit it was ambiguous as fuck add 50%IPrOH, 25%1- butanol, 25%GAA. No mention of nitroalkanes besides “below 20% excess”. I honestly hold all of the old members in extremely high regard but some things do not lend themselves to repeatability. I’m not trying to mock him at ALL but I think all scientists published or not need to ensure their methods are repeatable and easily understood hence moles and shit like that. While here in the US our research stations crumble around the fume hoods for lack of funding

Yes, I admit to running like ten experiments at once, and not being especially diligent with documenting them. I enjoyed visualizing the progression of this reaction using different solvent systems. And I certainly left out a lot of information in my writeup.

One thing I like about this reaction is that it can be run under a wide variety of solvent systems. Another thing I like about it is that you can tell how the reaction is progressing by monitoring the color change of the solution. Nitrostyrenes are typically yellow, and when a lot are present, the solution tends to look orange. Side products of this reaction are reddish-brown, and seem to form in the presence of water. Since the initial step involves an imine formation and the ensuing production of a water molecule, I started with about 2% molar mass (compared to benzaldehyde) of PEA free base. I figured this would be low enough that most of the water would be long gone (via azeotropic evaporation) before the reaction would progress to the point where it could interfere. If you're attempting to speed this reaction up though (I wasn't), then perhaps you could add PEA in portions, separated by a few hours. The goal is to make sure that the quantity of catalyst is low enough so that water concentration is as low as possible at any given time. By opting for a slow reaction, you also give your crystals plenty of time to grow, and thus you can get pretty big crystals this way.

However, because I was removing water via azeotropic evaporation, it seemed to matter a lot that I occasionally replace solvents. The shape of the container also seemed to be important, and ones with narrow openings obviously lost solvent at a lower rate. The reaction that I ran with n-butanol, GAA, and IPA had the most impressive results, but there seemed to be a complex system of azeotropes that I wasn't sure what to make of. IPA seemed to evaporate the most quickly, so I had to add new IPA more often than the other two. IPA is critical to this method, because it's responsible for the low solubility of the product, which will then precipitate crystals and remove itself from the reaction.

Of course, if you're removing water via azeotropic evaporation, you have to worry about the reactants evaporating too. The only one that evaporated significantly was nitromethane. So if you start with no excess of nitromethane, then there will not be enough of it for the reaction. But if you start with too much, then it will significantly dissolve the formed nitrostyrene, preventing it from forming huge crystals.

So, to conclude, I was vague was because I'd been topping off the solvents as they evaporated, since my goal was to combine the benefits of GAA (becomes acidic in the presence of water, stopping the reaction) and of IPA (high solubility of reactants, low solubility of product) and of n-butanol (forms a useful azeotrope with water).

[Edited on 7/20/18 by Melgar]




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[*] posted on 19-7-2018 at 11:35


Quote: Originally posted by Gl3n  
I had attempted to utilize Melgars ratios for large single crystals. I will admit it was ambiguous as fuck add 50%IPrOH, 25%1- butanol, 25%GAA. No mention of nitroalkanes besides “below 20% excess”. I honestly hold all of the old members in extremely high regard but some things do not lend themselves to repeatability. I’m not trying to mock him at ALL but I think all scientists published or not need to ensure their methods are repeatable and easily understood hence moles and shit like that. While here in the US our research stations crumble around the fume hoods for lack of funding


I think that if you look at his post not as how to do things exactly but as a soup of what might be possible. I took his post, not as a how-to, but more as to what are the possibilities. Much more a post to request a discussion rather than to state that he has all the answers.

This could be a good conversation. I personally would use well-vetted literature to guide me in an effort to do a somewhat similar synthesis.

This is a subject which does interest me, sometimes discussions are just that and if lucky a rare jewel is discovered.
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[*] posted on 20-7-2018 at 07:14


Quote: Originally posted by morganbw  
I think that if you look at his post not as how to do things exactly but as a soup of what might be possible. I took his post, not as a how-to, but more as to what are the possibilities. Much more a post to request a discussion rather than to state that he has all the answers.

This could be a good conversation. I personally would use well-vetted literature to guide me in an effort to do a somewhat similar synthesis.

This is a subject which does interest me, sometimes discussions are just that and if lucky a rare jewel is discovered.

Yeah, that was how I'd intended it. The photos and the description of the reaction conditions I ran was more of a proof-of-concept than anything, indicating that this could be a very fruitful direction for future research. However, if someone wanted to do this in a more systematic way that's more conducive to replication, I'd recommend starting by researching what azeotropes form from a mixture of water, isopropanol, GAA, and nitromethane. Then see if a different azeotrope forms when n-butanol is added to the mixture. Once it's established what azeotrope is removing water, then the solvent could be pre-mixed at that ratio, and added to the reaction flask as it evaporated. I would imagine that some nitromethane should be added to the pre-mixed solvent too, to account for the slow loss of it from the system. My theory is that the azeotrope that removes water will dominate when water is present, but some other azeotrope that removes nitromethane also manifests at times. Isopropanol, for instance, has a very significant azeotrope with nitromethane, which is something I learned only recently.

edit: Also, for anyone curious as to the reaction mechanism, the base that actually does the catalyzing is almost certainly the imine, NOT the amine. Imines are actually fairly basic in their own right, and certainly basic enough to catalyze this reaction. So if you're worried about making sure that there's enough free amine to catalyze this reaction, don't be. This is actually why methylamine works fairly well as a catalyst despite having a very low boiling point: it's almost never present as a free amine.

[Edited on 7/20/18 by Melgar]




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[*] posted on 28-7-2018 at 07:03


When I've done this reaction before with i-PrOH/GAA as the solvent I have always used a moisture trap. Nothing to remove the water created, but just to prevent more from absorbing into the solution since it's (initially) anhydrous.

If you do this in an open vessel over a matter of days wouldn't that become an issue? I would think at only 40-50C it would absorb atmospheric moisture faster than it drives it off as an azeotrope.
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[*] posted on 7-9-2018 at 10:41


Quote: Originally posted by Corrosive Joeseph  
Zinc-HCl reduces aromatic nitro and aliphatic terminal nitro groups to amines very well.

Secondary nitroalkenes give a mix of products which hydrolyze in excess acid to the ketone.

Don't forget this - http://orgsyn.org/demo.aspx?prep=cv1p0413

Attached here is by far the best review available on the net for nitrostyrene/nitroalkane reduction.

Apologies for slight topic drift. I will whip myself later.


/CJ


Has anyone here run the reaction as described in this paper (low temperature without any catalytic amine)? It seems superior to the traditional conditions of elevated temperatures and presence of a catalyst, in terms of simplicity and yield, but I can't seem many other cases of people carrying it out like this.
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Corrosive Joeseph
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[*] posted on 14-9-2018 at 02:36


Quote: Originally posted by monolithic  
Has anyone here run the reaction as described in this paper (low temperature without any catalytic amine)? It seems superior to the traditional conditions of elevated temperatures and presence of a catalyst, in terms of simplicity and yield, but I can't seem many other cases of people carrying it out like this.



OrgSyn procedures are usually rock-solid. The 'low temperature without catalytic amine' works for nitromethane but not for longer chain nitroalkanes. Just don't forget this -

"5. The alkaline solution must be added slowly to the acid, for the reverse procedure always forms an oil containing a saturated nitro alcohol. A large excess of acid at room temperature is used, conditions which facilitate the formation of the desired unsaturated nitro compound."


Oh BTW, I have heard great things regarding ethanolamine/acetic acid catalyst @ RT, based on the attached paper.

/CJ


Attachment: alizadeh2010.pdf (1MB)
This file has been downloaded 444 times





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[*] posted on 14-9-2018 at 06:15


Quote: Originally posted by Corrosive Joeseph  
Quote: Originally posted by monolithic  
Has anyone here run the reaction as described in this paper (low temperature without any catalytic amine)? It seems superior to the traditional conditions of elevated temperatures and presence of a catalyst, in terms of simplicity and yield, but I can't seem many other cases of people carrying it out like this.



OrgSyn procedures are usually rock-solid. The 'low temperature without catalytic amine' works for nitromethane but not for longer chain nitroalkanes. Just don't forget this -

"5. The alkaline solution must be added slowly to the acid, for the reverse procedure always forms an oil containing a saturated nitro alcohol. A large excess of acid at room temperature is used, conditions which facilitate the formation of the desired unsaturated nitro compound."


Oh BTW, I have heard great things regarding ethanolamine/acetic acid catalyst @ RT, based on the attached paper.

/CJ


Awesome find. Thank you for your contributions. :)
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[*] posted on 29-10-2018 at 08:52


Quote: Originally posted by alking  
When I've done this reaction before with i-PrOH/GAA as the solvent I have always used a moisture trap. Nothing to remove the water created, but just to prevent more from absorbing into the solution since it's (initially) anhydrous.

If you do this in an open vessel over a matter of days wouldn't that become an issue? I would think at only 40-50C it would absorb atmospheric moisture faster than it drives it off as an azeotrope.

Nope. The way azeotropes work is by having a lower boiling point than either of the constituent compounds. That also translates to having a higher vapor pressure, when the solvents are below the azeotropic boiling point. So any water that entered into the system would raise the vapor pressure, increasing the evaporation rate until it was gone, at which point the system would go back to having the lower vapor pressure consistent with an anhydrous solvent system.

The other key part here, is that you have to make sure that the solvent is warmer than the environment, so water can't condense on its surface. Also, that the evaporation rate is high enough to remove the water as an azeotrope as it forms.

So as long as solvent is evaporating at a positive rate, and the solvent is warmer than its surroundings, that shouldn't be a problem.




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[*] posted on 13-11-2018 at 15:31


Quote: Originally posted by Melgar  

Nope. The way azeotropes work is by having a lower boiling point than either of the constituent compounds. That also translates to having a higher vapor pressure, when the solvents are below the azeotropic boiling point. So any water that entered into the system would raise the vapor pressure, increasing the evaporation rate until it was gone, at which point the system would go back to having the lower vapor pressure consistent with an anhydrous solvent system.

The other key part here, is that you have to make sure that the solvent is warmer than the environment, so water can't condense on its surface. Also, that the evaporation rate is high enough to remove the water as an azeotrope as it forms.

So as long as solvent is evaporating at a positive rate, and the solvent is warmer than its surroundings, that shouldn't be a problem.


This is not correct. An azeotrope can also be higher boiling than either of its constituents. And an azeotrope (or atleast the values you find online) is measured at mixture reflux. At room temperature or slightly above these values would be completely different or even be zeotropic.
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[*] posted on 18-11-2018 at 09:48


Quote: Originally posted by Cactuar  
This is not correct. An azeotrope can also be higher boiling than either of its constituents. And an azeotrope (or atleast the values you find online) is measured at mixture reflux. At room temperature or slightly above these values would be completely different or even be zeotropic.

Let's just assume I specified "low-boiling azeotropes", since that's what I meant, and move on to the next part, ok?

My understanding of azeotropic evaporation is that what often happens is that the azeotropic solvent mixture cools to significantly below ambient temperature, due to heat loss from evaporation. This can induce water to condense on its surface, which changes the solvent composition, and throws off the azeotropic evaporation rate. Therefore, it should be possible (if this is correct, I could easily be wrong) to prevent this from happening by simultaneously limiting the exposed surface area, and ensuring that the solvent mixture is always ~5°C+ above ambient temperature.

I'm not exactly sure where this notion came from, but I'd appreciate it if someone could confirm a) whether it's definitively true or false b) if this phenomenon happens with azeotropic solvent mixtures that are hydrophobic or do not form azeotropes with water and c) if this is known to happen under an inert, anhydrous atmosphere. A quick literature search seems to indicate that evaporation rates of azeotropes are related in some capacity to their boiling-point azeotropes, though I'm not seeing much that's definitive on the subject.

Experimentally, 95% ethanol was left in two vessels, until half the original volume was gone. One vessel was a flat metal motion-picture film canister that was left at ambient temperature. The other was a much narrower still-picture film canister that was placed on the transformer for a router power supply, about 5°C above room temperature (room temperature fluctuated between 60°F and 66°F). When half of the alcohol was gone from each vessel, an attempt was made to ignite a piece of cotton dipped in each one. (The narrow vessel took longer for half its volume to evaporate, neither was timed) Both ignited, although the warmer one in the narrower vessel ignited noticeably more easily. Both were allowed to evaporate until only a small amount of liquid was left. The narrow, heated vessel ignited easily, but the liquid from the flat, unheated vessel would not ignite. There did not seem to be any drop in concentration in the heated alcohol in the narrow vessel.

Apologies for not having much in the way of equipment and/or instrumentation. This seems to at least demonstrate that water can be absorbed from the atmosphere such that it has an effect on azeotropic evaporation, if nothing else.




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[*] posted on 23-11-2018 at 14:24


Quote: Originally posted by Melgar  
My understanding of azeotropic evaporation is that what often happens is that the azeotropic solvent mixture cools to significantly below ambient temperature, due to heat loss from evaporation. This can induce water to condense on its surface, which changes the solvent composition, and throws off the azeotropic evaporation rate. Therefore, it should be possible (if this is correct, I could easily be wrong) to prevent this from happening by simultaneously limiting the exposed surface area, and ensuring that the solvent mixture is always ~5°C+ above ambient temperature.


I'm not sure about temperature dropping significantly below RT. I guess it depends on many things like the ambient temperature, the latent heat needed for vaporization, the area exposed and the vapor pressure of the liquid. If you drop acetone, DCM or ether on a watch glass this is clearly the case. However I don't think isopropanol in a beaker would cool down enough to make water start condensing on the walls and greatly change the solvent composition. It does seems logical though that a warmer mixture would have slightly less water in it, although any dry solvent, warm or not, would get wetted when exposed to the atmosphere.

Even if the mixture was zeotropic it wouldn't really matter in an open beaker, water would still evaporate along with isopropanol. I guess it would reach some sort of equilibrium where there is an equal chance of a water molecule being absorbed and one being evaporated. I don't know how they analyze compositions of azeotropes but I'm guessing it has to pass some fractionating column first.
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