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Borohydride synth?

Protium - 3-1-2005 at 00:49

How can one go about producing Sodium Borohydride?

Can this be done relatively easily in a makeshift lab with basic glassware?

JohnWW - 3-1-2005 at 01:59

The first step would be to get diborane, B2H6. (BH3 is electron-deficient, which is made up for in the dimer by means of three-center bonding). This is usually obtained by reacting elemental boron (or B2O3 may also be usable) with Mg, as mixed fine powders, in the absence of air (e.g. under argon) to form Mg3B2, which is hydrolyzed by dilute acids to Mg++ and gaseous B2H6, which has to be kept away from air. (If B2O3 is used, more Mg is required, and the product contains MgO which requires more acid to neutralize). The B2H6 is reacted with sodium hydride, NaH (from burning Na in H2 gas), in a suitable unreactive organic solvent, in the absence of air, to give NaBH4.

mick - 3-1-2005 at 09:09

You could try reacting lithium with and alkyl halide under dry conditions to form the alkyl lithium. React that in large excess with dry boric oxide, under dry conditions, and you might end up with a lithium tetra boroalkyl reagent. Might be as good as sodium borohydride but not as good as lithium aluminium hydride.

mick

solo - 3-1-2005 at 10:08

Here is a nice post with some patents for the synthesis of sodium borohydride.......solo

Attachment: NaBH4 patents.html (75kB)
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solo - 3-1-2005 at 10:20

The Preprationof Sodium Borohydride by the High Temperature Reaction of Sodium Hydride with Borate Esters
H.I.Sxhlesinger, herbert C. brown and A.E. Finholt
JACS 1975,205

http://home.ripway.com/2004-11/211899/JACS_75_205-1.djvu

Note: if not able to download wait 24 hrs then download, they only allow me 10mb /day of downloads

HRH_Prince_Charles - 3-1-2005 at 10:23

This paper from Rhodium reports no success with electrolytic reduction of borates to borohydride. The authors followed several of the patents with no success.

Attachment: borohydride.electrosynth.pdf (105kB)
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Marvin - 3-1-2005 at 10:37

Worth mentioning the 'standard' prep for many years, does not use the rather toxic and difficult to make diborane, trimethyl borate is a little on the toxic side but is almost trivial to make. Triethyl borate is much less toxic, only slightly harder to make and can also be used.

4 NaH + B(OCH3)3 => NaBH4 + 3NaOCH3

USP 2,461,661
USP 2,534,533
USP 2,683,721

It still leaves the problem of sodium hydride though.

JohnWW - 3-1-2005 at 11:00

What sort of environment, or solvent(s), are needed for that reaction?

BTW that home.ripway.com website is not operating.

Ref: Needed route to borates

solo - 3-1-2005 at 11:50

Procedures for Preparation of Methyl Borate
H.I.Schlesinger, Herbert C. Brown, Darwin L. Mayfield and James r. Gilbreath
JACS 1975, 213

Attachment: JACS_75_213-1.djvu (69kB)
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Ref: Sodium Borohydride Synthesis

solo - 3-1-2005 at 11:52

The available synthesis of sodium borohydride....now for those that can't download from ripaway....solo

Attachment: JACS_75_205-1.djvu (137kB)
This file has been downloaded 1733 times

Figured this was as good a place as any to post.

BromicAcid - 1-5-2005 at 11:41

Today I attempted to make MgB2 from magnesium shavings and boric oxide. 11 grams of boric oxide (probably not toally the oxide, probably some boric acid mixed in, dehydrated it a few days ago) and 15 grams of magnesium were jammed into a metal crucible in layers (as many of you know magnesium shavings take up quite a bit of volume so this whole mix was fairly volumous). The mixture was put into an isolated system (i.e., a one way gas valve was attached to the top) and after 45 minutes of heating with a propane torch, nothing. I removed it and found that the boric oxide with acid mixed in had semi melted and turned it into a maleable mass. This might be a good method to make ignots of this to run this reaction, i.e., melt boric acid and add enough magnesium shavings to it and stir.

Anyway, I had to heat the solid ball with a blow torch for about 3 mintues and then a bright light started that I continued to heat until it spread through the mass, after allowing to cool I found that half of the mass didn't react, the bottom half. So I removed it and heated that. What I ended up with was very similar to charcoal in color and how it felt when I crushed it. I ground it up noting a very few pieces of unreacted magnesium but overall it looks like a good reaction. Must find if appreciable MgB2 was formed.

A note on magnesium boride, I've found a number of references to it as MgB2 (magnesium diboride) but I've also found a few as Mg3B2. Not sure which is correct but since I found more references for the diboride I decided to shoot for that with my reaction:

B2O3 + 4Mg ---> 3MgO + MgB2

However what would the reaction for the magnesium boride reacting with HCl look like? I'm assuming its somewhat complicated being that diborane will decompose in contact with water, but with Mg3B2 the reaction can at least be made to appear simple:

Mg3B2 + 6HCl ---> B2H6 + 3MgCl2

BTW, my intention is not to use diborane to produce the borohydride but directly in reductions, it is more powerful from what I've read, able to reduce carboxylate groups whereas borohydride cannot.

BromicAcid - 5-5-2005 at 19:32

See the apparatus picture (attached) before reading the experimental section.

Today I decided I should test the potency of my magnesium boride. I took the product from my above reaction, 14 g and placed it into the bottom of the vessel shown in the picture filled with smoke. The vessel had a sepretory funnel in the top filled with HCl to drive the reaction foreward, out of the side of the two neck flask was a thermometer adaptor, the connection with the glass tubing teflon taped to prevent reaction. This glass tubing lead to the inlet of a gas bubbler and they were connected with a pice of rubber tubing, teflon taped on both sides, the glasses of each touching within. 100 ml of water was put into the gas bubbler and the vessel was evacutaed with a hand vacuum pump to decrease the explosion liklihood.

Upon adding HCl the flask immediately filled with smoke. (Fine B2O3 from reaction of diborane with the oxygen present?) The smoke carried over a little and bubbled through the bubbler. Noticeable magnesium shavings (though low in quantity) were still in the mixture so gas was expected from the reaction. I continued adding HCl as I have read that diborane hydrolyzes readily to boric oxide or acid in water (the intended goal being diborane produced would hydrolyze in the gas bubbler and be precipitated with HCl then filtered and weighed). I added a total of 100 ml of 28% HCl to the mixture, and lots and lots of gas came over, it had a unique odor, like cured ham mixed with gym socks.

When it was all over no boric acid precipitate was present and the water smelled of that strange smell. I added some HCl to the water in the bubbler hoping to precipitate some boric acid, nothing. Aside from the vessel filled with smoke the whole time and the strange smell nothing to show that diborane may have been produced.

It may be a better plan to use the product from the thermite reaction directly and mix with aquous NaOH under reflux to yield NaBH4 such as in the patents that are availible through the Hive material posted via Solo. Don't know about the solubility of NaBH4 but if it is fairly soluble then filtration and crystalization is the key being that the other things should be fairly insoluble, the only real impurity then would be borates.

smokefilled.jpg - 73kB

12AX7 - 5-5-2005 at 21:41

That smell is odd to describe, isn't it?

I burned some B2O3 + Mg/Al a while ago, it's been soaking in HCl since. Still some clumps, suppose I should filter and pulverize it some time...

Tim

Off Topic and Rambling

BromicAcid - 11-5-2005 at 19:05

No diborane can be detected by the reaction of magnesium boride and hydrochloric acid. After some investigation into the matter it was found that diborane hydrolyzes so readily that it only has fleeting existence when attempts are made to produce it in this way. The reaction with magnesium boride with HCl produces small amounts of hexaborane and others which were condensed by a number of chemists and fractionally distilled which yielded diborane. The yields are very low <1% in many cases but can be increased to as much as 11% by using 85% phosphoric acid. The same can be said of the reaction of sodium borohydride with hydrochloric acid which gives nearly quantitiative amounts of hydrogen, however with 85% phosphoric acid 40 - 50% yields of diborane can be had, higher concentrations give better yields but foaming and such can be a problem.

Magnesium diboride was prepared in US patent 2942935 by the reaction of elemental magnesium with boron at 950 C under an argon atmosphere (According to another paper if the temp is too high yields are reduced a thin iron or clay crucible should be used, nickel is attacked). The product was ~85% diborane with the remainder magnesium. Upon heating this with potassium hydroxide solutions, the higher the concentration the better. An 8 M KOH solution (250 ml) heated with 46 g MgB2 gave 12.4g potassium borohydride which was recovered by filtering, removing the water in vacuum, the KBH4 being the first thing to precipitate out. The reaction can also be carried out in alcohols or even in the solid phase. The following reaction is given in the patent for the formation of borohydrides assuming B2H6 is an intermediate product:

2B2H6 + 4MOH ---> 3MBH4 + MBO2 + 2H2O

In addition to treating potassium borohydride with phosphoric acid to give diborane, iodine, BF3 in ether, or mercurous chloride can accomplish this reaction.

Although solutions of THF*BH3 are common, diborane is 10x more soluble in dimethyl sulfide and looses less reactivity over time in it.

The reaction of diborane with Na/K alloy gives a salt that can be sublimed at 400C which has incredible reducing powder. A test for boranes is to bubble them though a solution of silver nitrate which causes a brown precipitate.

Sources:

Complete Treatise on Inorganic and Theoretical Chemistry pps 25-36

Recent Developments in the chemistry of Boron Hydrides H.I. Schlesinger ; Anton B. Burg 1941

Reduction of Organic Compounds with Diborane Clinton F. Lane Aldrich-Boranes, Inc. 1975

Nugget of information, might help

no1uno - 2-4-2009 at 09:25

The attached paper deals with the preparation of alkali metal borodeuterides from alkali metal alkoxides and trimethylamine.borodeuteride(d3).

Kinda nice, especially if we can access NaOEt from NaOH & EtOH. All we need now is a useful preparation of diborane that we can pass into the trimethylamine (would triethylamine be as effective?)...

Attachment: syntheses.alkali.metal.borodeuterides.pdf (969kB)
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Here is about the best idea I can see: http://pubs.acs.org/doi/abs/10.1021/ja01097a050

What about reducing with aluminium? Yes, it is useless for the reduction of borate to boron, they are combined, but that warning that I see every time I search for aluminium boride, suggests they give off borane(s) upon hydrolysis. Just a thought

[Edited on 2-4-2009 by no1uno]

sparkgap - 2-4-2009 at 09:39

Not too coincidentally, I was having yet another look at H.C. Brown's papers, so...

sparky (~_~)

Attachment: arcdibor.pdf (132kB)
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no1uno - 2-4-2009 at 17:00

Now, I'd have to wonder, is that pyridine borane which they purified by running the mixture of gasses through a dry-ice cold trap, comparable to the triethylamine.borane that they used in the Canadian paper? Wonder if we could use pyridine.borane to produce the sodium/lithium borohydride(s) or perhaps use trimethylamine/triethylamine to collect the relevant borane from the synthesis instead of pyridine?

Fuck, this looks tasty

PS BCl3 can be produced by the reaction of borates & PVC when they are burnt according to a study done by one of the fire brigade type journals. Yield is shit, but hell

Could one use aluminium diboride with dry HCl gas?

Attached is a file SOLO uploaded on another site.... Thanks & props are due to SOLO for it

PPS The separation of the mixtures, I wonder, could that be affected as H.C. Brown did it? Cold trap, then as the non-volatile triamine.borane?

Attachment: hydrides_of_boron._i._an_efficient_new_method_of_preparing_diborane-new_reactions_for_preparing_bromo-diborane_and_the_s (843kB)
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un0me2 - 28-7-2010 at 21:41

There is also this paper(Grab it while it's working).

PS I want this on top of the pile so I can print the other pages via another PC

zed - 28-7-2010 at 23:35

Ummmm. Most of the hot new work on Hydrides, is being done in the field of energy storage. Hydrides are formed, as a means of hydrogen production, for fuel.

These hydrides are usually formed by ball-milling. In the case of NaBH4, first Mg Hydride is formed by ball milling Mg in the presence of hydrogen. Then, the resultant MgH2 is ball milled with NaBO2 to produce NaBH4.

Unfortunately I cannot access the full paper on line. It is in the clutches of the American Chemical Society, and they want thirty bucks or so, for a full read.

Mechanochemical Synthesis of Sodium Borohydride by Recycling Sodium Metaborate

Lingyan Kong, Xinyu Cui, Huazi Jin, Jie Wu, Hao Du and Tianying Xiong*
State Key Laboratory of Corrosion and Protection, Institute of Metal Research, Chinese Academy of Sciences, Shenyang 110016, China
Energy Fuels, 2009, 23 (10), pp 5049–5054
DOI: 10.1021/ef900619y
Publication Date (Web): September 3, 2009
Copyright © 2009 American Chemical Society
*To whom correspondence should be addressed. Telephone: +86-24-23971746. Fax: +86-24-23971746. E-mail: tyxiong@imr.ac.cn.
Abstract

Hydrogen can be easily produced from the catalytic hydrolysis of sodium borohydride (NaBH4), with sodium metaborate (NaBO2) being the co-product. If NaBO2 can be economically recycled in the process, NaBH4 could be considered as a promising hydrogen carrier in fuel cells because of its high hydrogen content. In this paper, we report our investigation into the synthesis of NaBH4 from NaBO2 via ball milling. The starting materials for the synthesis were NaBO2 and magnesium hydride (MgH2). After ball milling at ambient temperature in inert gas, NaBH4 and magnesium oxide (MgO) were produced. The synthesized NaBH4 was extracted from the admixture by isopropylamine. Our results indicated that the yield of NaBH4 from this process was 71 wt % when the MgH2/NaBO2 mole ratio, ball/powder ratio (BPR), inert gas (argon) pressure, and milling time were 2.07:1, 50:1, 200 kPa, and 2 h, respectively. The mole ratio of the reactants (MgH2 and NaBO2), ball-milling time, BPR, and milling atmosphere were found to have significant influence over NaBH4 synthesis. We present the results and discuss the effects of these key parameters in this paper.

http://pubs.acs.org/doi/abs/10.1021/ef900619y

European patent for a similar synthesis of KBH4......

https://data.epo.org/publication-server/pdf-document?PN=EP15...

These works are based on an older US patent.

http://www.google.com/patents?id=0t9uAAAAEBAJ&printsec=a...

The most recent patent I came across, was assigned to Rohm and Haas. Sorry, no experimental. But, they claim good results by ball milling an active hydride, with NaBO2, in the presence of a solvent, at fairly normal pressures and temperatures. This document appears identical in most aspects to the European patent entered earlier. But, it has an expanded reference section.

http://www.freepatentsonline.com/7297316.html

Other material:

Kojima et al., Recycling Process of Sodium Metaborate to Sodium Borohydride, International Journal of Hydrogen Energy, vol. 28, pp. 989-993 (2003), no month.

Preparation of potassium borohydride by a mechano-chemical reaction of saline hydrides with dehydrated borate through ball milling

http://www.sciencedirect.com/science?_ob=ArticleURL&_udi...

Sorry, I don't have access to the full article.

[Edited on 29-7-2010 by zed]

stygian - 29-7-2010 at 18:51

awhile back I stumbled upon some interesting 'electrolytic borohydride' documents. Not the widespread "here is the science, and here is reality--doesnt work" documents, but ones talking about using polar aprotics as the medium. I think it also discussed alanates as well. I really wish I could find them again ;\ .. wheat from the chaff, as they say.

Sedit - 29-7-2010 at 19:02

Quote:

PS BCl3 can be produced by the reaction of borates & PVC when they are burnt according to a study done by one of the fire brigade type journals. Yield is shit, but hell Could one use aluminium diboride with dry HCl gas?

What about dehydration with H2SO4 or possibly Oleum. Fire in many cases causes a strong dehydration reaction and I wounder if this effect could be simular to the effects of strong acids on borates and PVC. You don't by any chance know where you heard this do you?

zed - 29-7-2010 at 19:11

Well, the drawback of the above Sodium Borohydride synthesis, is that you need an active metal hydride to make it with.

The good point, is that Calcium Hydride will suffice, and it is readily available commercially. It can also be made with some effort from Calcium Carbide.

Magnesium Hydride can also be self manufactured, and the metal is still readily available. Ball milling Magnesium under Hydrogen is one option, while another is producing the hydride via "Activated Magnesium".

Hoveland - 30-7-2010 at 14:46

Diborane can be made by fusing sodium boride with NaOH.
6NaOH + 4B --> 2Na3BO3 + B2H6
White Phosphorous reacts similarly
3NaOH + P --> Na3PO3 + PH3
"(phoshine) can be made by boiling white phosphorus in a solution of potassium hydroxide" http://mysite.du.edu/~jcalvert/phys/phosphor.htm
However the yield of diborane from this method is very low, because the diborane has a tendancy to react back with the sodium hydroxide.
6NaOH + B2H6 --> 2Na3BO3 + 6H2
Using a flux of potassium oxide might help, since the formation of K3B would not be favorable. Also, I think I remember that metallic sodium was involved in the mixture too, but I do not think this is absolutely necessary.
Further reading ,under "Boron decombustion" toward the end
http://www.eagle.ca/~gcowan/boron_blast.html

Just to clarify, both diborane and BH4(-) react with water.
Another interesting fact, although I do not have the reference, diborane reacts with anhydrous ammonia, mainly forming (NH3)2BH2(+) BH4(-), whereas an adduct of BH3 with ether reacts with NH3 to form only NH3BH3.

[Edited on 30-7-2010 by Hoveland]

un0me2 - 31-7-2010 at 05:12

Diborane "can" be made by reducing gaseous Boron Halides (see the 2nd paper).

Boron Halides can be made via the reaction of the relevant Aluminium Halide - ie. AlCl3 with BF3 (for the preparation of BF3 see the 3rd paper, and for the preparation of Boron Halides from there, see papers 4-6). That is why I am so interested in the low-temperature chlorination of Aluminium thread, if that works we have a real shot.

Borohydrides can be made with diborane and the relevant alkali metal alkoxide (I wonder if we could use the glycoxide? _- for the preparation see the paper from Hoekstra #6), although it also forms an adduct with THF which is supposed to be a useful reducing agent in its own right (so too the amine.boranes). There are several reviews attached too.

As you will see with the preparation of the fluoroborate - the ammonium fluoride + H2SO4 sold at the local car accessory shop/hardware store for cleaning wheel rims should be OK. Drying the boric acid (from the pool store would appear to be unnecessary), the fluoroborate salt(s) seem like they are easier to dry than the boric acid (which is a bitch to dry). The reaction between the BF3(g) and the relevant Aluminium Halide looks simple enough, although obviously ensuring everything is dry as possible would be essential.

The use of a glass tube with a bed of aluminium filings/powder to reduce the Boron Halide to Diborane @350C looks well within the limits of ordinary Borosilicate glass, massive amounts of cooling would be essential, but I'm sure one or more of the contributor's here can work out how to harness the Diborane

What is the maximum temperature at which it could be used to make an aminoborane/adduct with a solvent?

Attachment: Burkhardt.Matos.Boron.Reagents.in.Process.Chemistry.Excellent.Tools.for.Selective.Reductions.pdf (829kB)
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Attachment: Hurd.The.Preparation.of.Boron.Hydrides.by.the.Reduction.of.Boron.Halides.pdf (455kB)
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Attachment: Schlessinger.Inorganic.Laboratory.Preparations.1962.Preparation.Fluoroborate.pdf (95kB)
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Attachment: Vorobyova.Practical.Inorganic.Chemistry.Boron.Halides.from.Aluminium.Halides.pdf (92kB)
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Attachment: Gamble.Gilmont.Stiff.The.Reaction.of.Boron.Trifluoride.with.Aluminium.Chloride.or.Bromide.pdf (359kB)
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Attachment: Hurd.The.Synthesis.of.Boron.Trichloride.pdf (245kB)
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Attachment: Hoekstra.The.Preparation.and.Properties.of.Alkali.Metal.Borohydrides.pdf (1.1MB)
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Attachment: Brown.Rao.Hydroboration.III.The.Reduction.of.Organic.Compounds.by.Diborane.An.Acid.Type.Reducing.Agent.pdf (773kB)
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[Edited on 31-7-2010 by un0me2]

un0me2 - 31-7-2010 at 06:59

And, if you have a nice, inert atmosphere to do it in, Lithium Triethylborohydride (aka super hydride, stronger than LiAlH4 apparently) is made via the reaction of Triethylboron with Lithium Hydride in THF. According to a paper just found, it is fucking near OTC (with the minor issue that triethylboron burns hot as hell immediately it comes into contact with air, on the plus side AlCl3 is reportedly the by-product in 95% yield).

Ok, considering how wonderful the compound seems to look like to work with, both preparations I can see in the last two attached papers claim 100% yields from triethylboron and LiH in THF.

Standard Solution of Lithium Triethylborohydride

In a dry 500-mL flask fitted with a side arm, a rubber syringe cap, and a magnetic stirring bar, 7.2 g (900 mmol, 50% excess) of lithium hydride was placed, and a reflux condenser connected to a mercury bubbler was attached. After 314.8 mL of THF was introduced, the system was flushed with nitrogen. While the mixture was vigorously stirred, 85.2 mL (600 mmol, total volume of the solution 400 mL) of triethylborane was introduced slowly. Generally, an exothermic reaction begins some 5-15 min following the addition. At this stage, an ice bath was placed under the flask to control the reaction and to avoid overflow of the reaction mixture through the condenser. After this vigorous reaction was over, the reaction mixture was refluxed for 2-3 h in order to ensure completion. The resulting solution was filtered through a filter chamber fitted with a sintered-glass (fme-grade) fiiter under slight positive pressure of nitrogen in order to remove excess lithium hydride. The resulting clear solution was standardized by removing an aliquot, hydrolyzing it with a water-glycerine-THF (1:l:l) mixture, and measuring the hydrogen evolved. With a series of preparations, the concentrations were determined to be in the range of 1.45-1.55 M in LiEt3BH. The THF solution of lithium triethylborohydride is characterized by a strong, broad absorption in the IR at 2060 cm-l (BH)3.” If the solution was maintained under a dry nitrogen atmosphere, no change in composition was detected in months at room temperature or in days at 65 “C (refluxing THF).

Attachment: Zhao.etal.A.Novel.Method.for.the.Preparation.of.Diborane.from.Aluminium.Ethyl.Chloride.Boron.Oxide.and.Hydrogen.pdf (346kB)
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Attachment: Brown.Super.Hydrides.pdf (1018kB)
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Attachment: Brown.Kim.Krishnamurthy.Selective.Reductions.26.Lithium.Triethylborohydride.as.an.Exceptionally.Powerful.and.Selective.R (1.6MB)
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Attachment: Krishnamurthy.Brown.Selective.Reductions.31.Lithium.TriethylBorohdyride.as.an.Exceptionally.Powerful.Nucleophile.pdf (892kB)
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[Edited on 1-8-2010 by un0me2]

benzylchloride1 - 31-7-2010 at 21:35

The procedure for synthesizing sodium hydride from sodium metal can be found in Inorganic Syntheses Collective Volume V on pages 6-13. The procedure details the production of an extremely fine dispersion of sodium metal and the subsequent reaction with hydrogen between 220 C -280 C at atmospheric pressure in standard laboratory glassware. This could be of use if sodium metal is available since sodium hydride can be reacted with methyl borate to form sodium borohydride.

un0me2 - 1-8-2010 at 14:40

Actually at STP (or anywhere near it) the only product found by Schlesinger/Brown was Sodium Trimethoxyborohydride, the reaction of that in THF with Diborane is pure sodium borohydride though. At high temperature and pressure, NaH does react to form "some" borohydride. I'm waiting on the reference, but that is what the first page says.

The file attached "isn't" the one I was thinking of, but it does show the use of solvent for the preparation of Lithium Borohydride, Sodium Trimethoxyborohydride and from THAT in solvent with diborane, one can achieve Sodium Borohydride.

[Edited on 2-8-2010 by un0me2]

Attachment: Brown.Tierney.The.Reaction.of.Lewis.Acids.of.Boron.with.Sodium.Hydride.and.Borohydride.pdf (947kB)
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un0me2 - 3-8-2010 at 23:57

Here is me looking at triethylborane and looking to see what is the minimum level of oxygen with which it reacts, only to find these articles (here, here & here) in which they researched the radical reactions of triethylborane-air @ 20C. Fair enough, they did only open the things up after the reaction had started (and everything was in solution), but that would have took balls, BIG BALLS

- they also cite this paper where 600mL was good for 16 restarts/afterburner lighting on the SR-71 Blackbirds Combined Operation Turbojet/Ramjet-type engines.

Radical reactions? What like people running like fuck when they realise what the clown in the next booth is going to do?

[Edited on 4-8-2010 by un0me2]

Chainhit222 - 4-8-2010 at 13:11

Quote: Originally posted by benzylchloride1

The procedure for synthesizing sodium hydride from sodium metal can be found in Inorganic Syntheses Collective Volume V on pages 6-13. The procedure details the production of an extremely fine dispersion of sodium metal and the subsequent reaction with hydrogen between 220 C -280 C at atmospheric pressure in standard laboratory glassware. This could be of use if sodium metal is available since sodium hydride can be reacted with methyl borate to form sodium borohydride.

here you go:

http://img231.imageshack.us/img231/6498/sodium1.jpg
http://img571.imageshack.us/img571/8318/sodium2.png
http://img155.imageshack.us/img155/3707/sodium3.jpg

http://img641.imageshack.us/img641/7985/nah1.jpg
http://img375.imageshack.us/img375/3885/nah2.jpg
http://img641.imageshack.us/img641/5019/nah3t.jpg
http://img52.imageshack.us/img52/2431/nah4.png
http://img8.imageshack.us/img8/30/nah5.jpg

[Edited on 4-8-2010 by Chainhit222]

un0me2 - 4-8-2010 at 18:51

It is probably easier for all if you just print it - use PrimoPDF as a printer, then just print the pages - to get the effect you see in the attached PDF's, one is the same preparation of Sodium Dispersions and the other the preparation of Lithium Hydride (in a glass tube with a Mecker burner @500C).

Now I'm wondering, has anyone seen the preparation of LiH in dispersion? I know I've seen papers on making Lithium Sand in order to make Alkyllithium reagents, much the same as the sodium dispersion is made. So it strikes me that it should be possible to prepare LiH in much the same way as is done with the sodium dispersion... Albeit at a higher temperature maybe, but surely it would react?

Attachment: Inorganic.Syntheses.Vol.5.pp.6.10.Sodium.Dispersions.and.Sodium.Hydride.pdf (274kB)
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Attachment: Vorobyova.Practical.Inorganic.Chemistry.1987.pp.184.5.Preparation.of.Lithium.Hydride.pdf (93kB)
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Chainhit222 - 4-8-2010 at 19:15

Quote: Originally posted by un0me2

It is probably easier for all if you just print it - use PrimoPDF as a printer, then just print the pages - to get the effect you see in the attached PDF's, one is the same preparation of Sodium Dispersions and the other the preparation of Lithium Hydride (in a glass tube with a Mecker burner @500C).

Now I'm wondering, has anyone seen the preparation of LiH in dispersion? I know I've seen papers on making Lithium Sand in order to make Alkyllithium reagents, much the same as the sodium dispersion is made. So it strikes me that it should be possible to prepare LiH in much the same way as is done with the sodium dispersion... Albeit at a higher temperature maybe, but surely it would react?

I bet you that it will work just for for LiH. You can use the hydrogen test described in the NaH procedure to see if its working (compare two bubblers, if output is bubbling slower then input then it will be turning into LiH)

[Edited on 5-8-2010 by Chainhit222]

un0me2 - 4-8-2010 at 20:15

Well the alternative is always the preparation then the thermal degradation of alkyl (ethyl) lithium to ethylene and LiH, which is about as short a route (wasteful though - 50% loss off the bat) as probably exists. Although doing so directly from ethyllithium, as in the attached paper (cited many, many places (such as here & having been reported by Schlesinger & Brown, I'll take them at their word - it ain't some dodgy patent).

The second paper (attached) catalytic degradation of Ethyllithium into LiH + ethylene (good only if you are mad enough to still want triethylborane and/or super hydride after reading the other paper).

[Edited on 5-8-2010 by un0me2]

Attachment: Schlesinger.Brown.Metallo.Borohydrides.III.Lithium.Borohydride.pdf (181kB)
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Attachment: Zgonnik.etal.Reactions.of.Organometallic.Compounds.with.HeavyMetal.Salts.4.Reaction.of.Ethyllithium.with.Titanium.Trichl (108kB)
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sodium borohydride from common reagents

slinky - 17-2-2011 at 07:04

CH4 + NaBO2 Heat -> NaBH4 + CO2

Sodium metaborate (borax can be used) is strongly heated in an atmosphere of methane (in home natural gas is 97% methane) which produces sodium borohydride and carbon dioxide. The suggested reaction temperature is 900C. I'm thinking I'll use a stainless steel vessel with an inlet tube extending to just above the borax melt. As the sodium borohydride is formed it will boil and I should be able to condense it with a coldfinger. The vessel will have a an exit port at the top for purging unreacted natural gas which will sweep away the CO2. The reaction vessel will look something like this:

I have searched the forum and do not see any other chatter about the aforementioned reaction. Has anyone here attempted this synthesis? If anyone has any details, references, or comments about this reaction I'd love to read them. Here's what I could dig up with patent searches and google.

Attachment: methane.borax.to.sodium.borohydride.us7019105.pdf (184kB)
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Attachment: methane.borax.to.sodium.borohydride.us7294323.pdf (72kB)
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Attachment: millennium.cell.inc.review.chemical.processes.pdf (404kB)
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Attachment: us.dept.of.energy.fuels.of.the.future.for.cars.and.trucks.pdf (502kB)
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wireshark - 2-2-2013 at 21:09

Quote: Originally posted by slinky

CH4 + NaBO2 Heat -> NaBH4 + CO2

I've skimmed the third document before. Look at it again. The ΔG using methane is too positive for this reaction to proceed "at any reasonable temperature."

If you can't get borohydride, use reductive hydrogenation. Often the point of using metal hydrides is convenience, so it's pretty ridiculous to try to make them. Anyway, metal hydrides are impractical for industry and the chemical-lacking chemist alike.

[Edited on 3-2-2013 by wireshark]

Magpie - 2-2-2013 at 21:26

Quote: Originally posted by wireshark

Anyway, metal hydrides are impractical for industry....

Not all industry. The pulp bleaching industry uses NaBH4, probably by the ton. See Wiki. Also see the very informative publication by Rohm & Haas linked in the Wiki article references.

Forum member ordenblitz thought he had stumbled on a facile method for the home chemist but he dropped his research. I tried to pick up where he left off but also had no success.

[Edited on 3-2-2013 by Magpie]

S.C. Wack - 3-2-2013 at 06:51

Quote: Originally posted by wireshark

I've skimmed the third document before.

And I did it just now, as it's one of few non-patent google hits mentioning the MgB2 -> KBH4 process of JACS 78, 4176 (1956) and any patents they wrote. People here are clearly on to this but I couldn't find any actual mention of the little article...http://pubs.acs.org/doi/abs/10.1021/ja01597a093

"A 13% conversion of boron to borohydride was obtained, as determined by the amount of hydrogen evolved upon acidification of the solution."

simba - 4-2-2013 at 13:45

Quote: Originally posted by S.C. Wack

Quote: Originally posted by wireshark

I've skimmed the third document before.

13% sounds very good if this method really works, looks pretty easy, almost too easy, in fact.

Orenousername - 18-4-2016 at 22:03

Has anyone tried MgH2 + NaBO2 yet? and do the reagents need to be anhydrous?

zed - 24-4-2016 at 15:17

Sure, this has been done, but probably not by anyone here. And, yes....the reagents would have to be completely anhydrous.

Also do-able with CaH2, which was formerly an inexpensive and easy get. MgH2 works better though,

Check via the search engine. There is a lot of information on this topic.

[Edited on 24-4-2016 by zed]

CRUSTY - 12-6-2016 at 12:50

Quote: Originally posted by Orenousername

Has anyone tried MgH2 + NaBO2 yet? and do the reagents need to be anhydrous?

I would assume they need to be anhydrous, since NaBH₄ reacts with water. I would try this, but MgH₂ seems like a total pain in the ass to work with, as you'd need to run the hydrogenation (as well as the MgH₂ synthesis) under an inert, anhydrous atmosphere, and you have to some sort of milling setup and a very fine magnesium powder in order to get a reasonable yield, since H₂ gas can barely diffuse through MgH₂ IIRC.

clearly_not_atara - 12-6-2016 at 18:14

I don't think anyone has successfully produced the hydrides. Everyone focuses on what you do after you get the hydride because nobody can figure out how to make one lol.

It's known that magnesium in napthalene containing catalytic -- [TiCl4 or TiBr4] EDIT: CrCl3 works which is WAY better -- can be hydrogenated at 20 C at normal pressure.

http://onlinelibrary.wiley.com/doi/10.1002/anie.198008181/fu...

I think this is probably going to be the type of procedure that eventually succeeds. The ball milling procedure mentioned on page 1 is touchy and requires activated magnesium in a perfectly dry atmosphere. TiCl4/naphthalene acts as a powerful in situ dehydrating system to get the reaction going that's just more practical... it's possible that the requisite catalyst loading is small.

Bromine is dangerous, but the reaction of a tiny amount of bromine with titanium can't be too bad, right? :p

The trimethoxyaluminum hydride ion can then be made (I think) by reaction of a metal hydride with Al(OMe)3. The latter is made by simply reacting activated aluminum with methanol, i.e. Al + 3 MeOH + cat HgCl2 >> Al(OMe)3 + 3/2 H2. Magnesium trimethoxyaluminum hydride is the presumptive reaction product of MgH2 + Al(OMe)3.

https://en.wikipedia.org/wiki/Reductions_with_metal_alkoxyal...

EDIT: Actual paper attached. CrCl3 instead of TiCl4 is a *big* improvement. EDIT2: FeCl2 is even better.

Attachment: 10.1002@anie.198008181.pdf (246kB)
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[Edited on 13-6-2016 by clearly_not_atara]

zed - 13-6-2016 at 13:47

"Bromine is dangerous, but the reaction of a tiny amount of bromine with titanium can't be too bad, right? :p"

Bromine is dangerous, but we used to routinely use it in lower level chemistry courses. I can remember a classmate getting bromine all over his fingers. He got it off right away, and as I recall, there were no serious consequences.

I might be concerned about Bromine reacting with Titanium. Titanium is pretty non-reactive, but once it decides to react, it might not be easy to put the brakes on.

clearly_not_atara - 13-6-2016 at 14:23

The article also mentions that FeCl2 can be used as a catalyst. Assuming this is correct (and doesn't actually refer to FeCl3), anhydrous FeCl2 can be produced using methanolic HCl with vacuum drying:

https://en.wikipedia.org/wiki/Iron(II)_chloride#Laboratory_preparation

This is probably the easiest way to make one of the transition-metal catalysts. I don't think it was a typo, either -- FeCl3 seems too oxidizing to be used in the preparation of a hydride.

NB: the produced MgH2 is pyrophoric!

[Edited on 13-6-2016 by clearly_not_atara]

gdflp - 13-6-2016 at 16:21

For anyone interested, here is the article referenced by the wiki page :

Attachment: FerrousHalides.pdf (37kB)
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Corrosive Joeseph - 4-12-2017 at 19:33

Found this and thought I'd drop it here............

"Nanocrystalline magnesium for hydrogen storage"

https://doi.org/10.1016/S0925-8388(99)00073-0

Abstract

The hydrogen storage properties of MgH2 are significantly enhanced by a proper engineering of the microstructure and surface. Magnesium powders are produced in a nanocrystalline form, which gives remarkable improvement of absorption/desorption kinetics. Ball milling, which is used for fabrication of nanocrystalline magnesium, improves both the morphology of the powders and the surface activity for hydrogenation. The hydriding properties are further enhanced by catalysis through nano-particles of Pd located on magnesium surface. Nanocrystalline magnesium with such a catalyst exhibits an outstanding hydrogenation performance: very fast kinetics, operation at lower temperatures than conventional magnesium and no need for activation.

And still using the ball mill............

"Preparation of sodium borohydride by the reaction of MgH2 with dehydrated borax through ball milling at room temperature"

https://doi.org/10.1016/S0925-8388(02)00872-1

Abstract

A convenient method was developed to synthesize NaBH4 by the reaction of MgH2 with Na2B4O7 through ball milling at room temperature. In order to improve the sodium borohydride yield, Na compounds were added to compensate the Na insufficiency in reactants when MgH2 instead of NaH was used as the reducing agent. It was found that Na2CO3 addition was better than NaOH or Na2O2 addition in increasing the borohydride yield.

/CJ

[EDIT] - Attachment
[EDIT2] - Attachment2
[Edited on 5-12-2017 by Corrosive Joeseph]

Attachment: li2003.pdf (551kB)
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[Edited on 5-12-2017 by Corrosive Joeseph]

Attachment: zaluska1999.pdf (630kB)
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clearly_not_atara - 9-11-2023 at 10:35

Quote: Originally posted by S.C. Wack

Quote: Originally posted by wireshark

I've skimmed the third document before.

We should contextualize this. I believe the balanced reaction should be:

4 KOH + 2 MgB2 >> 2 MgO + KBO2 + KBH4 + 2 B

So the ideal yield is only 25% and 13% molar yield is 52% of theory.

The remaining boron can be recovered and reused quite feasibly if the solution is acidified since we should see precipitation of boric acid and of course elemental B is not soluble.

The last question is how to optimize the reaction:

Mg + B + B2O3 >> MgO + MgB2

to obtain a reaction which is self-sustaining but not too exothermic and makes good use of the Mg. I suspect that the product mixture can be separated by using cold aqueous acetic acid solution to dissolve MgO without much decomposition of MgB2.

[Edited on 9-11-2023 by clearly_not_atara]