Electra - 13-6-2014 at 10:06
Not sure if this belongs here or in reagents....
I got my hands on some Sodium Borohydride, wanted to do some hydrogenations to play around with, but don't want to waste it. I also have a number of
different Phase Transfer Catalysts. I found this link:
http://phasetransfercatalysis.com/ptc_tip/use-ptc-to-save-a-...
This isn't the first source I've read about using borohydride with a PTC. Large amounts of excess are not required, you can use almost quantitative
amounts without any going to waste apparently. Only problem I am facing is that wikipedia (bad source I know) says that Sodium Borohydride reacts with
water. I am going to assume it does not react to a large extent if 97% yields can be achieved in an aq. medium with only 4% excess based on the 4
hydrides.
I see that the authors of the above article use 1M of NaOH and state Borohydride is stable under the right PH in water. Can anyone else testify to
this or comment on its stability?
An additional question is, what would its solubility be, since Wiki just says "soluble, reacts with water"? Am I to assume high?
[Edited on 13-6-2014 by Electra]
Metacelsus - 13-6-2014 at 10:16
In my organic lab class, we ran a borohydride reduction of a ketone in methanol. Borohydride does react with methanol, but it was slow enough that it
didn't matter much. Also, we were using a 1.5x excess of borohydride.
I assume that the rate of hydrolysis in water is comparable to the rate in other protic solvents.
Anyway, don't try it with LAH, and you'll be fine. 
forgottenpassword - 13-6-2014 at 10:31
There's lots of information about solubilities; decomposition rates and pH of aqueous solutions, etc. here: http://www.dow.com/assets/attachments/industry/pharma_medica...
[Edited on 13-6-2014 by forgottenpassword]
Electra - 13-6-2014 at 11:20
Great info guys. I am curious on the specific role NaOH plays in stabilizing the Borohydride, in terms of mechanism.
forgottenpassword - 13-6-2014 at 12:22
BH4- reacts with H+. Adding NaOH simply reduces the amount of H+. pH is intrinsically linked to the concentration of H+ in a solution, of course.
[Edited on 13-6-2014 by forgottenpassword]
Electra - 13-6-2014 at 14:43
Perhaps you can help me understand something.
Are all four hydrogen active as hydrides in sodium borohydride? If so, what is the byproduct of a typical hydrogenation?
I ask because the above article had me thinking...
http://phasetransfercatalysis.com/ptc_tip/use-ptc-to-save-a-...
They show the ketone was hydrogenated with 1/4 molar equivalent of borohydride in very high yields. How is this possible? Even if all hydrogen were
used as hydrides, each ketone still needs at least two hydrogen to become an alcohol. Meaning a minimum of 1/2 mole eq to get a 97.7% yield that they
claim to have gotten.
I saw another article that implies that only 1 hydrogen from each borohydride is used, leaving BH3 as a product, when ketones or aldehydes are
hydrogenated.
http://www.masterorganicchemistry.com/2011/08/12/reagent-fri...
In any case, what are the byproducts I would be looking at for a ketone/aldehyde hydrogenation? Something insoluble? Or a soluble salt?
edit:
Judging by
http://www.chemgapedia.de/vsengine/vlu/vsc/en/ch/2/vlu/oxida...
It seems all four hydride are used and B(OH)3 - boric acid is formed as a result, but I don't see where the extra hydrogen come from. Even this link
says one mole of ketone is reduced by 1/4 mole of borohydride. The stoichiometry doesn't work out for that. To turn 1 mole of ketone to 1 mole of
alcohol you need 2 total moles of hydrogen for each ketone. Using a 1/4 equiv of Borohydride only provides half of that.
Unless the extra hydrogen comes from water? If that were the case there would still be 1 hydrogen short.
[Edited on 13-6-2014 by Electra]
forgottenpassword - 13-6-2014 at 15:20
Yes, all of the four hydrides are active in water. If you look here: http://masterorganicchemistry.com/wp-content/uploads/2011/08...
You will see that BH3 and OH- are generated from the first reaction. These combine to give HOBH3- which is itself a hydride donor. Via a similar
mechanism you get (HO)2BH2- and (HO)3BH- respectivrly from the two successive hydride transfers. When the final hydride is lost, boric acid is the
final product.
[Edited on 13-6-2014 by forgottenpassword]
zed - 25-6-2014 at 16:48
True, NaBH4 breaks down in methanol. It slowly fizzes away hydrogen. I'm not sure how significant the loss actually is. I always assumed the
hydrolysis-breakdown product to be Sodium Borate.
In my hydrogenation experiments, I always used excess NaBH4 to compensate for losses.
This was a problem, as NaBH4 was expensive at the time.
Better to run your hydrogenation experiments with NaBH4 dissolved in ethanol. Sadly, the use of pure ethanol comes with its own problems, due to its
tax status. And, many denatured types of ethanol are unsuitable for use with NaBH4 due to the presence of ketones and esters, or methanol.
FireLion3 - 27-6-2014 at 01:40
Electra I'm trying to do the same thing. I found a chart somewhere that shows the stability of sodium borohydride in different PH's.
Dr.Bob - 27-6-2014 at 05:58
The NaBH4 removed H+ ions when dissolved in water, thus shifting the pH, just like the NaOH, only NaOH is cheaper. That is why once you have added
enough NaBH4, the reaction with water slows, and why NaOH slows the reaction.
But you will loose some hydride in water, no matter what. But since each BH4 generates up to 4 hydrides, you can get away with not too much BH4.
Good luck.