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Harmine, harmaline and THH from Syrian Rue. Verification and finetuning of the VDS-protocols Options
 
Jees
#21 Posted : 10/21/2016 3:44:32 PM

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ijahdan wrote:
...Just one thing, when chemists talk about 'excess' base, does this just mean more than is needed to reach the maximum pH? How much more?
Max pH? That is getting hard to nail so up high pH levels. Therefore I gave the volume numbers: I began with 100 ml of a 7% vinegar + 3 gr zinc and needed 60 ml of a 8mol/L base (160 gr dry lye in 0.5 liter water) to make all zinc hydroxyde turn into liquid again, meaning go trough a filter while THH should be caught.

The excess-base-trick is proven by now, but actually one we would like to skip because we don't know how THH would stand such a treat.
For now, we hope to not use excess base but instead pH gentler ways ammonia or sodium carbonate for obvious reasons. The results of that should hopefully make part of this thread later on, so that we can acknowledge for ourselves the document/paper in the OP.

Will make a dilute base and try to dissolve the zinc hydroxide in it, fingers crossed.
We're not done yet Pleased
 

Trippy glass for trippy people.
 
An1cca
#22 Posted : 10/21/2016 3:46:13 PM

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Hey ijahdan,

I just completed some verification experiments on Van Der Sypt's protocols. I'll post the results in a few hours. For the moment, let's say that I miss an 'ecstatic' emoticonThumbs up
 
Jees
#23 Posted : 10/21/2016 4:20:34 PM

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Jees wrote:
...Will make a dilute base and try to dissolve the zinc hydroxide in it...

According to the paper page8 point 2.5, all THH was precipitated at pH 9.8
So I think to make the dilute base minimal but just reaching that value.
This handy site says for sodium carbonate that 10 milligrams (or 0.1 mmol/L) does that exactly in 1 liter water.
 
An1cca
#24 Posted : 10/21/2016 7:46:31 PM

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Hey Jees, in the protocols at the end of the article, VDS washes the alkaloids with ammonia 3% or Na2CO3 0,5%. Presumably this is what he refers to as 'dilute base'.

And a handy site indeed, thanks!
 
An1cca
#25 Posted : 10/21/2016 8:06:37 PM

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So this is what I did:

Added 6g of Zn-powder to 200ml of acetic acid 7%. Let it react for 21 hours. A constant production of hydrogen was observed. Filtered off zinc with dry weight of 4,938g. So more or less 1g of Zn reacted. Unreacted Zn can be reused, so this seems very economical.

Then I divided the 200ml of solution and to half of it, I added 100ml demineralized water (to arrive at the same concentration as in the VDS-protocol after washing). To this solution, strong ammonia 12% was added in small increments. Between pH 5 and 9 (crudely), a fine mist appeared in the solution (pic 3). However, at pH 10: abracadabra: the solution turns perfectly clear (pic4)! This is exactly what we hoped for and what is described in the online chemistry course that Pitubo referred to (attached the relevant screenshot, thanks Pitubo!).

To the other half I also added 100ml of demineralized water, but then added concentrated sodium carbonate solution (20g/100ml) in small increments. Again, starting more or less at pH5, a fine precipitate appears. However, unlike this fine mist as obtained from ammonia, in thise case the 'fluffy' precipitate tends to agglomerate into small floating clusters (clouding) that have a tendency to slowly settle (pic9). Even after adding excess base (18g), the precipitate remained. At this moment, I'm trying to filter the solution and it looks like a pain in the Crying or very sad. I'll keep you posted about the residue's solubility...
An1cca attached the following image(s):
Chemguide hexaaquazinc.jpg (51kb) downloaded 338 time(s).
pic1 Before ammonia.jpg (71kb) downloaded 336 time(s).
pic2 At pH 7,13.jpg (70kb) downloaded 335 time(s).
pic3 fine 'mist'.jpg (53kb) downloaded 334 time(s).
pic4 At pH 10,12.jpg (78kb) downloaded 334 time(s).
pic5 foaming with sodium carbonate.jpg (66kb) downloaded 332 time(s).
pic6 At pH 5,93.jpg (69kb) downloaded 332 time(s).
pic7 'flocculent' precipitate.jpg (54kb) downloaded 331 time(s).
pic8 At pH 7.jpg (69kb) downloaded 331 time(s).
pic9 At pH 10,51 corrected white balance.jpg (44kb) downloaded 331 time(s).
 
An1cca
#26 Posted : 10/21/2016 8:36:16 PM

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Quote from wikipedia 'Zinc hydroxide':

"If excess sodium hydroxide is added, the precipitate of zinc hydroxide will dissolve, forming a colorless solution of zincate ion:

Zn(OH)2 + 2 OH− → Zn(OH)42−.
This property can be used as a test for zinc ions in solution, but it is not exclusive, since aluminum and lead compounds behave in a very similar manner. Unlike the hydroxides of aluminum and lead, zinc hydroxide also dissolves in excess aqueous ammonia to form a colorless, water-soluble ammine complex.

Zinc hydroxide will dissolve because the ion is normally surrounded by water ligands; when excess sodium hydroxide is added to the solution the hydroxide ions will reduce the complex to a −2 charge and make it soluble. When excess ammonia is added, it sets up an equilibrium which provides hydroxide ions; the formation of hydroxide ions causes a similar reaction as sodium hydroxide and creates a +2 charged complex with a co-ordination number of 4 with the ammonia ligands - this makes the complex soluble so that it dissolves."

Currently, I'm drying my precipitate to know its dry weight. Then I'll try to dissolve it in ammonia 3% and Na2CO3 0,5% as stated in the VDS-protocols.
 
Jees
#27 Posted : 10/21/2016 9:41:06 PM

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Nice!

And I had something similar going on today with ammonia, could not form precipitation Shocked
But must re check everything, yet I love to hear what you report Love this looks promising.

* * *

The story of the dilute base (sodcarb & also tried out ammonia) to dissolve already dry ZnOH makes one looking a bit stupid Laughing
Maybe it works different for ZnOH that has not been dried, I dunno, but these fails today aren't very promising. Noticed that dried ZnOH took it's fair time to dissolve in citric water pH 2.5. Makes me no wonder that the dilute-base-trick didn't party.
 
An1cca
#28 Posted : 10/21/2016 10:06:21 PM

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Well, this gel-like precipitate has a hard time drying.

In order to determine its solubility, I decided to take a few pinches of the residue and drop them into:

-ammonia 3%: sparingly soluble
-ammonia 12%: sparinlgy soluble (a bit better?)
-sodium carbonate 0,5%: insoluble
-vinegar: gives off carbon dioxide

So the residue proves to be zinc carbonate.

Then, when I added some ammonia 12% to the filtrate that nevertheless had still some clouding in it, the mist instantly vanished. This confirms the -albeit- low solubility in ammonia. However, it seems to me that it is too low to be readily washed out from between alkaloids.

Perhaps this is the reason why VDS only used ammonia in this experiment? In experiment 2.7, sodium bicarbonate is used first. The precipitate that ensues weighs about 2 times as much as expected. I suppose the non-harmine fraction is zinc carbonate. Then he continues to get rid of the salt by doing a Manske or an ammonia precipitation from the dissolved acetate.

My conclusion is that if available, ammonia should be used in the basification of the reactants. If only sodium(bi)carbonate is available, a Manske is necessary to clean up the residue. However, as far as I get it, in the case of experiment 2.7, the use of sodium bicarbonate is obligatory because its use bypasses the need for pH-metry in a combined harmine-THH mixture. However I think he might have made a mistake by stating that the harmine-fraction can be cleaned by dissolving it in vinegar and reprecipitating with Na2CO3. Will zinc carbonate not again precipitate out?

Another conclusion can be that a crude post-Manske harmine-harmaline Rue alkaloid extract can be converted to harmine-THH without using sodium bicarbonate but by using pH-metry to differentiate between the 2 alkaloids while precipitating. In this case, far less unwanted zinc-salts will co-precipitate (only the 'mist'-kind of precipitate seen above, that might be readily removed by washing?).

This information is exciting! The use of ammonium chloride appears to be superfluous. If VDS's protocol now also proves to work while effectively synthesizing THH, it would well open doors for many people being reluctant to use hydrochloric acid, lye, ammonium chloride or organic solvents. And he claims a yield of 83%!

I'll start working on the separation method described in the paper (protocol 2.1) and keep you posted...

An1cca attached the following image(s):
Precipitate from excess carbonate.jpg (48kb) downloaded 323 time(s).
 
An1cca
#29 Posted : 10/21/2016 10:36:32 PM

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I think I forgot another option and research question: if sodium bicarbonate is in fact able to precipitate all zinc ions as zinc carbonate and if it is unable to precipitate THH, then another ammonia-free option might be

1) 3g harmaline + 100ml acetic acid 7% + 3g zinc
2) let react for many (12?) hours
3) filter off zinc and recycle
4) add large excess of sodium bicarbonate (can even be solid)
5) filter (taking out excess sodium bicarbonate and presumably all zinc carbonate)
6) add excess of concentrated sodium carbonate solution (20g/100ml) to filtrate
7) filter off THH and wash with dilute sodium carbonate solution (0,5%?)

If I can obtain some pure harmaline by VDS's selective precipitation-protocol, I will test this synthesis route.

I think I'll take a break nowCool
Happy researching!

 
endlessness
#30 Posted : 10/21/2016 10:39:20 PM

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I love this researching spirit, keep it up, family Cool Thumbs up
 
Jees
#31 Posted : 10/22/2016 12:16:55 AM

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An1cca wrote:
I think I forgot another option and research question: if sodium bicarbonate is in fact able to precipitate all zinc ions as zinc carbonate and if it is unable to precipitate THH, then another ammonia-free option might be...
There's a tremendous two if's in that sentence, so it may be, but prove is out there.

But if you don't suffer from one or other ammonia phobia then there's this:
we had a few indicators here already that ammonia keeps the zinc hydroxide from precipitating without the use of ammonium chloride. Time to dig into that.
However the dilute-base-ammonia-wash could not dissolve dry zinc hydroxide, the question at hand is what happens on precipitated zinc hydroxyde in solution.

I took around 250 mg dry zinc hydroxide and dissolved it in kitchen 7% acetic acid, I think it was like 25 ml. Some heat and time and all dissolved. This state simulates a reduction step after which the excess zinc was removed, only now there are no alks in because we want to see what happens on the zinc hydroxide front solely.

I wanted to be sure to really see all that zinc hydroxide as precipitate flocculant swimming. So I added sodium carbonate until pH 9.5 and I suppose most or all of the zinc hydroxide indeed precipitated. Now the magic: adding 1 ml of 12%ammonia and all was rocket fast dissolved, clear liquid all the way, and end pH was 10.

So with the zinc hydroxide in suspension (not dry) the astonishing little 1 ml ammonia does the trick just as well as a whopping 40 ml of 8mol lye.
Shocked
And to know that we now end at pH10 instead of 14, I think THH would be grateful.

So we have magic to make zinc hydroxide liquid at pH10, while we suppose in this configuration the THH would not follow the ZnOH's example. Then THH could be filtered off.

After this experiment things start to give hope.
I will repeat the experiment using only ammonia as a base right after the excess zinc removal step, then pH right up to 10. It would be mimicking the OP paper but with no alks present. It would allow us to see what happens strictly on the zinc level solely.

ToBeContinued

 
Jees
#32 Posted : 10/22/2016 12:36:45 AM

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Jees wrote:
... Now the magic: adding 1 ml of 12%ammonia and all was rocket fast dissolved, clear liquid all the way, and end pH was 10...
Now this is funny:
so the ZnOH dissolved but how to recover it from here? Can't filter it out at this pH10 Laughing

Answer: get rid of the 1 ml ammonia. So I put the vial in a warm water bath and as the ammonia started to vape off, the zinc hydroxide began to solidify again.
Life is wonderful Big grin

Add again 1 ml of 12% ammonia, all clear again.
What more do we want?

Sell your ammonium chloride shares now Twisted Evil
 
pitubo
#33 Posted : 10/22/2016 2:05:20 AM

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I have a hunch that the ammonium chloride acts as extra HCl.

2 (NH4)+ + (Zn(H2O)6)2+ --> (Zn(NH3)2(H2O)4)2+ + 2 H+

Just a wild guess.. I am not a trained chemist, really.
 
DreaMTripper
#34 Posted : 10/22/2016 5:20:47 AM

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Interesting thread! Will get my pencil and paper out later see if I notice anything different in the reactions along the way. That clear solution is a beautiful sight by the way!
You don't need any comprehensive training in chemistry for balancing equations just need to know the basics and the likely products (see solubility rules) but looks like everyone has a good handle on that. One thing is that if there is to be a precipitate then the charges must be balanced. They can easily be found online. Another thing to take into account is the stoichiometric ratios, once you have discovered this by balancing the equation then you can calculate how much Zn is needed to fully react by using the mole to mass relationship. That is mass = moles x formula mass. moles = mass/formula mass. Molarity (concentration) = moles/volume and moles of a solute = Molarity x Volume (litres) .
It may have been mentioned before apologies if I missed it but what is the state of THH in an acidic solution? Would it be possible to reduce pH with minimal acetic acid add Zinc to reduce to THH then add sodium carbonate (or bicarbonate so pH doesn't go too high?) to precipitate Zinc carbonate (and possibly any other contaminate, cadmium is a nasty chemical), filter the this out then raise the pH to 9+ with NaOH to precipitate the THH while simultaneously creating soluble sodium zincate out of any unreacted zinc as an extra separation step to be safe. After precipitation and separation you could even acidify the solution, add sodium carbonate again ,filter it then weigh it, then again with the previously precipitated zinc carbonate to determine if all the Zn has been removed. That requires knowing what percentage of the compound Zn is then using that to judge what percentage of the final precipitate is Zn.
 
An1cca
#35 Posted : 10/22/2016 7:51:42 AM

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Jees wrote:
But if you don't suffer from one or other ammonia phobia then there's this...


According to my psychiatrist, it has something to do with phallic shaped ammonia bottles in my childhood Big grin


Jees wrote:
I will repeat the experiment using only ammonia as a base right after the excess zinc removal step, then pH right up to 10. It would be mimicking the OP paper but with no alks present. It would allow us to see what happens strictly on the zinc level solely.


I think this what I've done already in the above experiment, where the first half of the reacted solution was treated with ammonia 12%. Actually, I added 37ml of it, strictly adhering to VDS's protocol. Or do you mean something else?

Currently, I'm finishing the separation experiment. I will post the results shortly...


@DreaMTripper: It's best to use an excess Zinc to make sure this will not be the limiting factor. And it now looks like lye could be avoided altogether.
 
Intezam
#36 Posted : 10/22/2016 8:24:13 AM

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Great job An1cca, Jees Thumbs up Love Cool

DreaMTripper wrote:
(....and possibly any other contaminate, cadmium is a nasty chemical)


Cadmium? Is it a common (pyrotechnical) zinc powder contaminant? Is there reason to be worried?
 
pitubo
#37 Posted : 10/22/2016 11:38:11 AM

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Intezam wrote:
DreaMTripper wrote:
(....and possibly any other contaminate, cadmium is a nasty chemical)

Cadmium? Is it a common (pyrotechnical) zinc powder contaminant? Is there reason to be worried?

Cadmium, lead, arsenic and other minor contaminants. So yes, avoid pyrotechnical zinc. You need at least Super High Grade zinc, but preferrably even more pure. When buying zinc dust, always ask for an analysis certificate.

https://en.wikipedia.org/wiki/Zinc_refining
https://en.wikipedia.org/wiki/Zinc_smelting
 
Jees
#38 Posted : 10/22/2016 11:51:14 AM

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An1cca wrote:
...Currently, I'm drying my precipitate to know its dry weight. Then I'll try to dissolve it in ammonia 3% and Na2CO3 0,5% as stated in the VDS-protocols.
I think we should leave the path of trying to nail down the Zn(OH)2 properties when that has been taken out of the solution. Concerning the transitions between dissolved-precipitated, for now I suspect the ZnOH to behave differently when it is within or out of the solution.
Since the aim is to eventually do everything in a liquid state, maybe it's better to do all our dissolving attempts in such a liquid state, not on dry ZnOH.

I see use for dried ZnOH only to make a weight statement (as you/we were doing), or to put it into vinegar later to dissolve and have it from there to experiment further with.

* * *

I agree with the excess zinc style, all that has not been reacted into ZnOH is reusable so far as myself and An1cca could notice. Decanting your liquid leaves most zinc at the bottom to re use. Some fine zinc mist will likely be in the decanted liquid, so that has to be cotton ball filtered before base is added.
After washing any vinegar of the zinc to keep, I put a tiny water level on the reusable zinc to have a molecular seal from the air, we don't the want zinc oxidized.

* * *

An1cca wrote:
...I think this what I've done already... Or do you mean something else?...
True, no real need to do same over again. Thumbs up

I'd like to change to this question: is it possible to precipitate all ZnOH with sodium bicarbonate? Have the solids filtered out, then base the filtrate further (with sodium carbonate or lye but sure no ammonia) to see if any more ZnOH goes solid.
This experiment could pave the way to the ammonia-free route you were telling about.

But I have my reservations about it all: after filtering out the ZnOH you were telling that the filtrate still had a mist in it (that could be vanished with some ammonia). Well I noticed the same thing, as if it was impossible to filter out all precipitated ZnOH, my filtrate too had significant residual "mist". The filtering is no good means here to rely on, it seems. Like once ZnOH is in precipitate state it is something elusive to catch with confidence.

This affects much what DreamTripper suggested too:
Quote:
... add sodium carbonate (or bicarbonate so pH doesn't go too high?) to precipitate Zinc carbonate (and possibly any other contaminate, cadmium is a nasty chemical), filter the this out then raise the pH to 9+ with NaOH to precipitate the THH...
as we both found it hard to filter out ZnOH with confidence.

For this reason maybe leaving all the routes that involve "catching the ZnOH with a filter" and go rather for tricks only that make ZnOH dissolved. What you think?

* * *

pitubo wrote:
I have a hunch that the ammonium chloride acts as extra HCl.
2 (NH4)+ + (Zn(H2O)6)2+ --> (Zn(NH3)2(H2O)4)2+ + 2 H+
Just a wild guess.. I am not a trained chemist, really.

Nor a chemist here, just found this link:
http://chemiday.com/en/r...n/3-1-0-10704?q=balance

So far things look promising to not need that trick after all, it seems extra ammonia on its own works too per the "ligand exchange action". Love that term Smile

 
An1cca
#39 Posted : 10/22/2016 11:54:00 AM

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So this is how did a preliminary check of another paragraph in the paper: the pH-selective separation of harmine and harmaline.

In line with Van Der Sypt's experiment 2.3, I dissolved 6g of post-Manske (3 times) harmine+harmaline freebase in 35ml acetic acid 7% + 250ml of demineralized water. A little heating and some stirring quickly succeeded in dissolving all of it. After cooling, I added ammonia 12% dropwise under continuous stirring and observation of the solution's pH.

While adding the ammonia, the pH took off from around pH 4, rose quickly around pH 5 (typical titration-curve?) and again, slowly climbed up to around 7. Around pH 6,5 every drop of ammonia caused a cloud of yellow precipitate in the dark-brown solution that progressively stayed longer. And then, around pH 7,15, a strange phenomenon occured: the more base I added, the lower the pH became! After adding a drop of ammonia, the pH quickly rose, but then steadily declined to around 6,5. This depression of the pH, that Van Der Sypt also noted, seems to know an exponential decrease. A few hours of waiting caused the pH to go all the way down to 6,3! Then I continued adding base and noticed that the speed of pH-depression/time got smaller, up to around pH 7,5 where even the smallest addition caused manifest and permanent rising of the pH. Here the solution was filtered again. On pic1 you see the fraction that was filtered at that time as well as the color of the mother liquor.

As predicted in the article, the pH rose to about 7,8 (without much precipitation) until suddenly, the solution turned cloudy again and the pH started to crash. Macroscopically, flickering appeared in the solution (harmaline plates reflecting light?). This fraction was filtered as well and the addition of ammonia was resumed. Pic2 shows the situation at pH 7,95 where pH-depression/time again started to diminish. Sadly, the flickering is difficult to photograph (for me that isEmbarrased ). Pic3 shows the fraction that was taken at that pH. Notice how the color of the mother-liquor turned lemon-yellow now. When more ammonia was added, the solution turned cloudy again, but precipitation proceeded slowly and not much harmaline was recovered after that. I'll keep you posted about the dry weights.

So my own 'basification curve' seems to be analogous to the one in the article. If the analysis of the reactants proves to be correct as well, we finally have an EXACT way of separating harmine and harmalineThumbs up . Specifically, for the study of harmaline reduction, our starting material should not be contaminated by harmine.

Now, I don't have access to TLC, melting point determination or GC/MS. Perhaps I can gain access to a microscope. Does anyone of the Nexians have the possibility to do a qualitative and quantitative analysis on this material?

As far as I'm concerned, I feel rather Big grin about his!
An1cca attached the following image(s):
pic1 nearing the pH-peak between harmine and harmaline.jpg (46kb) downloaded 261 time(s).
pic2 Precipitation at pH 7,95.jpg (95kb) downloaded 262 time(s).
pic3 Precipitate and mother liquor at pH 7,95.jpg (94kb) downloaded 260 time(s).
 
An1cca
#40 Posted : 10/22/2016 12:04:03 PM

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Jees wrote:
I think we should leave the path of trying to nail down the Zn(OH)2 properties when that has been taken out of the solution. Concerning the transitions between dissolved-precipitated, for now I suspect the ZnOH to behave differently when it is within or out of the solution.
Since the aim is to eventually do everything in a liquid state, maybe it's better to do all our dissolving attempts in such a liquid state, not on dry ZnOH.


I agree with that!


Jees wrote:
I'd like to change to this question: is it possible to precipitate all ZnOH with sodium bicarbonate? Have the solids filtered out, then base the filtrate further (with sodium carbonate or lye but sure no ammonia) to see if any more ZnOH goes solid.
This experiment could pave the way to the ammonia-free route you were telling about.


Good thinking, gonna check this out...



Jees wrote:
For this reason maybe leaving all the routes that involve "catching the ZnOH with a filter" and go rather for tricks only that make ZnOH dissolved. What you think?


Yep, I agree this would be easiest wey to go. Let me embrace my darkest fears and welcome ammonia!Wink Nice working with you, Jees!
 
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