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Author: Subject: Quantitative preparation of boric acid (from borax)
Fulmen
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[*] posted on 26-4-2020 at 01:18
Quantitative preparation of boric acid (from borax)


First off: No, I can't just buy it. It's simply not sold here.

But I did get hold of a few pounds of borax. And preparation is pretty simple, just add a stronger acid. But which one? I happened to have a few pounds of sodium bisulfate (pool ph-) I didn't really need, so I opted for this.

I'm working with batches of 150g borax, adding 120g bisulfate gives a pH of appr 3, which should be enough to ensure complete conversion.

That brings us to the real challenge, separation. The bisulfate (or sulfuric acid) is actually a decent choice, as the sodium sulfate has a slightly negative solubility from 35 to 100°C (490 to 425g/l). This should precipitate fairly pure boric acid (270-70g/l). But there is still boric acid left in solution, so first I collect the fraction from 35 to room temp and add it to the next batch.
I can then refrigerate the remaining solution to extract the sodium sulfate.
This will of course contain a fair amount of boric acid, so a second crystallization would increase yields and purity.

Question is, are there better acids for this? Hydrochloric acid would produce sodium chloride which has a fairly neutral solubility (366-390g/l), but the boric acid will probably be less pure, so a second recrystallisation will be required.




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[*] posted on 26-4-2020 at 02:18


Crystalise out the boric acid as far as you can.
Then dry the solution down to form crystalline sodium sulphate, sodium bisulphate, and boric acid.
Extract that with alcohol to recover most of the remaining boric acid.
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Fulmen
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[*] posted on 26-4-2020 at 02:52


Alcohol was on my list, but I had trouble finding solubility data on it. (Digression: I feel that Google has become noticeably worse in the last year or so...) I have some EtOH, but it's dyed so it would probably require further purification. And I don't like spending more money and chems than I have to, even though it's not a huge expense one should always try to keep cost and complexity down.

I'm also running out of bisulfate, if there are better acids now's the time to change method.




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[*] posted on 26-4-2020 at 03:00


Would Methanol or Isopropyl Alcohol also work for this purpose, I wonder?

Also, if you're wanting a change from Google, try DuckDuckGo or Startpage.com :cool:
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[*] posted on 26-4-2020 at 03:35


I think methanol would be a better bet than IPA.
Also, now that I think of it, it's probably better to have a slight excess of borax when you extract into alcohol, it might esterify under acid conditions.
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[*] posted on 26-4-2020 at 05:52


Quote: Originally posted by unionised  
I think methanol would be a better bet than IPA.


What makes you say that, as a matter of interest? Is it that Methanol and Ethanol are more similar in their properties as solvents?

[Edited on 26-4-2020 by G-Coupled]
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Fulmen
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[*] posted on 26-4-2020 at 07:24


MeOH should be the most polar alcohol, so it's not unreasonable to assume it's better for ionic compounds. However it's not readily available here. Also, without some solubility data it quickly becomes another research project. And I hadn't even considered esterification, I'll have to look into that first.

Part of the reason for asking is that I'm running low on bisulfate. So if there is a better option now's the time to change method.




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[*] posted on 26-4-2020 at 07:35


Honestly I just used HCl, it was quick and simple, then I just recrystallized in H2O. The solubility of boric acid in water at 0C is so low you won't lose much, and the purity will be quite high because NaCl has a high solubility. I did the same with potassium sorbate.
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[*] posted on 26-4-2020 at 12:28


yeah, I've come to roughly the same conclusion. The bisulfate was chosen mainly because I had a kilo I didn't need and I was out of HCl. But I think HCl would have been simpler. Oh well, we live and learn.



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[*] posted on 26-4-2020 at 13:26


Quote: Originally posted by Fulmen  
yeah, I've come to roughly the same conclusion. The bisulfate was chosen mainly because I had a kilo I didn't need and I was out of HCl. But I think HCl would have been simpler. Oh well, we live and learn.


This is a bit of a process but if you have a distillation set you could probably take NaCl and bisulfate to produce HCl vapour in much the same way bisulfate is used with NaNO3 to produce HNO3 and bubble it through a solution of sodium borate.
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