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happyfooddance
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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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alking
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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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Gl3n
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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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Gl3n
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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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Metacelsus
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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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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  |
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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morganbw
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| 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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Melgar
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| 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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alking
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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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monolithic
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| 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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| 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 554 times
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monolithic
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| 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
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Awesome find. Thank you for your contributions. 
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Melgar
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| 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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Cactuar
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| 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.
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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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Melgar
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| 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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Cactuar
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| 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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