Sciencemadness Discussion Board - Sodium borohydride conservation

Sodium borohydride conservation

Melgar - 27-6-2017 at 22:39

I figured that this would be the appropriate forum, since a gram of NaBH4 saved is a gram of NaBH4 earned, after all.

I noticed that certain procedures sometimes called for adding sodium borohydride to water or a water/alcohol solution, then waiting for the bubbling to subside before proceeding. Then I realized that the only reason for this was to raise the pH. Well, I had a reaction that called for sodium borohydride, and it didn't seem like a high pH would be detrimental, so instead of adding sodium borohydride, I added about a quarter the amount of sodium borohydride that it called for and let it dissolve. After adding the borohydride, there was virtually no bubbling even though a lot of it dissolved. Neat. Then I added the first portion of the second solution. There was a fine effervescence throughout the solution, and the flask got slightly warmer in my hand. The solution gradually changed color from yellowish-orange to clear. A tingle actually went down my spine because of how clean the reaction looked. Why doesn't anyone ever think to recommend stuff like this? Do they think we're MADE of sodium borohydride? It could also allow borohydride reactions in methanol, since one of the main problems with that solvent is the fact that sodium borohydride doesn't have a pH-raising effect on it like it does water. I haven't actually had a reason try this yet, but if anyone has the means to quantify the results of this type of reaction (methanol with pH adjusted using sodium hydroxide), I'd certainly be interested.

Dr.Bob - 28-6-2017 at 11:48

You could likely just add some sodium acetate, carbonate or hydroxide to the reaction first and then once the pH is raised a bit, then add the borohydride slowly. I think that you are right, as I have done the same things, and once the protons are all used up, the borohydride survives long enough to react. Also, many preps call for one equivalent of borohydride per ketone, and since BH4 should be able to provide up to 4 hydrides, that is often a vast excess, which you simply have to quench later during the workup.

The key is that borohydride is often considered cheap compared to the compound you are reducing, so little thought is given to the amount of borohydride. I have used some compounds where I know the prep uses too much of one reagent to try to drive it to completion, but often it is way more than needed if you titrate the reaction carefully. But that takes time and effort, which many industrial chemists don't want to waste.

Melgar - 28-6-2017 at 12:34

Quote:

But that takes time and effort, which many industrial chemists don't want to waste.

I take it commercially-available sodium borohydride doesn't require an icepick to get out of the jar, then?

I meant to say that I added some sodium hydroxide initially, then borohydride, and the reaction started out slow, rather than with lots of hydrogen bubbles.

Loptr - 29-6-2017 at 06:02

Quote: Originally posted by Melgar

Quote:

But that takes time and effort, which many industrial chemists don't want to waste.

I take it commercially-available sodium borohydride doesn't require an icepick to get out of the jar, then?

I meant to say that I added some sodium hydroxide initially, then borohydride, and the reaction started out slow, rather than with lots of hydrogen bubbles.

Commercial aqueous solutions of sodium borohydride are just that, as it's in equilibrium with hydroxide formation. I have a bottle of it around here somewhere (my hobbies have taken over).

EDIT: I now see that you are talking about Methanol+NaOH followed by addition of NaBH4, and not just an aqueous solution.

[Edited on 29-6-2017 by Loptr]

zed - 1-7-2017 at 16:35

Don't know how stable Na Borohydride is, in MeOH/NaOH. I have used Na Borohydride in straight Methanol, and I consider the results....not good.

Oh, I suppose it works....but the Borohydride clearly decomposes at a fairly rapidly rate. The Hydrogen just fizzes away! And, Dammit, I don't want to pay for that kinda side action!

Now, guys that obtain their Borohydride in industrial/bulk quantities, might be largely unconcerned by this money fizzing away. Sodium Borohydride is fairly inexpensive, if you buy a hundred pounds or so. Around 50 bucks a pound, last time I checked. About 10 grams fer a buck.

For those of us paying closer to 1 dollar per gram, the expense is maddening.

Now, if you are suggesting that adding a little base to your methanol solvent, might prevent Borohydride from prematurely decomposing........ I like the idea!

Using Ethanol as a solvent, is effective, but it presents other problems. Tax, expense, and possible legal entanglements.

Melgar - 1-7-2017 at 18:37

Quote: Originally posted by zed

Now, if you are suggesting that adding a little base to your methanol solvent, might prevent Borohydride from prematurely decomposing........ I like the idea

That IS what I'm suggesting! I actually just tested it myself in a couple of test tubes. One on the right had like 15 mg of NaOH added to it before the NaBH4. As you can see, a lot fewer bubbles. It definitely appears to work how you'd expect it to.