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Author: Subject: A/B extraction Raising pH question
szuko03
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[*] posted on 17-2-2016 at 07:23
A/B extraction Raising pH question


This seems like a stupid question to me but inorganic chemistry was never my strong suit. I have been attempting to isolate some alkaloids in plant material and have gotten to the step of adding a base to the aquious solution in an attempt to liberate it from the salt complex to dissolve it in an inorganic solvent, the whole standard procedure thing.

Anyway my main question is when the liquid was acidic, I used dilute acetic acid, the pH was about 2-3, now I have added an extreme abundance of sodium hydroxide. My digital pH meter has stopped at 10.5 - 10.75 and it does not appear to increase even though I have added more NaOH then I would have predicted I would need.

Is my assumption that after a specific point, when all of the acidic ions are neutralized by the NaOH, the pH will continue to rise correct? I would like to believe that the experiment is proceeding as it should but I do not wish to dump more lye into the solution only to have it be pointless as I wish to not waste money, materials or time.

If anyone can shed some light as to why the pH has held at 10.5 and wont increase I am all ears. I set out to do this to understand inorganic chemistry better and here I am trying :)

[Edited on 17-2-2016 by szuko03]




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[*] posted on 17-2-2016 at 09:01


Are you sure your pH meter is correct? They can sometimes be inaccurate at high pH (this phenomenon is called sodium error, and depends on the electrode used).

Edit: Sodium error would be an error in the opposite direction to what you observed, so it isn't likely the culprit. Still, I would suggest testing your pH meter under alkaline conditions.

How much excess NaOH did you use, exactly?

[Edited on 2-17-2016 by Cheddite Cheese]




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[*] posted on 17-2-2016 at 09:07


Extractions on biological matter has to deal with a high buffer capacity from all the miscellaneous crud that is also present. It's pretty common to need a much larger quantity of acid/base than someone(inexperienced) would expect to reach a desired pH.

Also, recalibrate your pH meter.
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szuko03
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[*] posted on 17-2-2016 at 09:22


Thats the thing though, it may not be "calibrated" as best it can be but our "hard" tap water reads 7.5-7.75. I know thats not much of a calibration but it at least speaks of its accuracy in my mind. But yes after noticing this i did plan on getting some calibration equipment for the meter.

Also the aqueous layer is about 2L and the total amount of lye added is about 1kg (dont judge if it seems like a lot trying to be safe) I suppose the simplest answer is the pH meter but I just dont get how my "hard" tap water readings could be so accurate, other then the alluded to idea that extremes are harder to gauge. It should also be added that this was done in purchased distilled water, not the tap water. I also did have minor success when I pulled the freebase form out with a nonpolar solvent and salted the desired compound out... minor being like 5-10% of the target amount. I know its there I figured the issue was the pH wasnt allowing a good percentage of the compound to be in its alkaline form.



[Edited on 17-2-2016 by szuko03]




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[*] posted on 17-2-2016 at 10:28


Why not get a smaller sample of the extract, 20 ml or so, and add a decent amount of concentrated aqueous NaOH to it and extrapolate from that?

It'd be far more convenient then groping around in the dark with 1L + mixtures
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[*] posted on 17-2-2016 at 10:32


Quote: Originally posted by semesa  
Why not get a smaller sample of the extract, 20 ml or so, and add a decent amount of concentrated aqueous NaOH to it and extrapolate from that?

It'd be far more convenient then groping around in the dark with 1L + mixtures


So true, I was just excited and its not like the materials cost is excessive in my mind. But yeah I do small scale for everything else so go figure lol.

Thanks for the advise though I am trying to see now if doing another nonpolar extraction on the basic aqueous layer will yield more so I can at least test if adding more sodium hydroxide pushed the equilibrium and I was able to isolate more, I figured that was the logical next step.




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[*] posted on 19-2-2016 at 11:02


Quote: Originally posted by semesa  
Extractions on biological matter has to deal with a high buffer capacity from all the miscellaneous crud that is also present. It's pretty common to need a much larger quantity of acid/base than someone(inexperienced) would expect to reach a desired pH.



Seconding this. I have an experience of extracting anthocyanin (a natural pH indicator) from hibiscus herbal tea. The leaves acted as a pretty powerful buffer and kept acidifying my solution as I kept adding ammonia to reach the neutral purple color. Only when I filtered the leaves off I was able to reach pH 7.




Smells like ammonia....
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