I am wondering of the possibility to make boron hydrides (boranes?) at home.
As on-hand chemicals I have sodium perborate, 95% H2SO4, H3BO4, NaOH, 30% NH4OH ecc.
Other chemicals required are not real a problem (depending on which ones), but
particular apparatus, yes this is my problem.
If anyone could post a synthesis of (if possible) liquid boron hydrides, or else other BnHn.
Thanks at all for help! Cyrus - 24-4-2005 at 15:50
Are you sure you don't mean H3BO3 instead of H3BO4? I dunno... but I also want to get into some more boron chem. experiments... sorry I
can't help. neutrino - 24-4-2005 at 16:48
Boric acid is H<sub>3</sub>BO<sub>3</sub>, or B(OH)<sub>3</sub>.
I don't have any experience in boron's chemistry, but I would guess that you'd want to start with simple diborane
(B<sub>2</sub>H<sub>6</sub>. To make this, you could
probably run a thermite reaction between boric oxide (made by heating the acid) to get a boride and adding an acid to make diborane. That’s my
guess, anyway.BromicAcid - 24-4-2005 at 17:15
I've posted this theoretical syntesis elsewhere but here it is again.
(Note that aluminum will not work as a substitute for magneium in this reaction as aluminum boride is inert and will
not react to form the hydride under normal conditions.)
3) Drip cold HCl on magnesium boride to give you a mixture of boric acid and magnesium chloride and other things in your reaction vessel, diborane and
higher boranes will come off the reaction vessel, they are easily hydrolyzed and as well all know, easily oxidized. Normally they would be lead into
something like THF to form a stable complex that could be easily used.
Now it stands that higher boranes are indeed liquid and higher still are polymetric solids. To attain these higher boranes pyrolysis is usually used
which leads to a mixuture of boranes from what was once mostly diborane. So alternating chambers of heat and cold (to condense out your liquid
boranes) might be a possibilty.
Boranes burn with a beautiful green flame, they create very fine dispersions of boric oxide which is not something considered good for your lungs and
they are incredibly toxic for some odd reason.Chris The Great - 10-10-2005 at 22:11
I'm curious as to whether the reaction of boron halides with alkali hydrides will yield diborane in decent yeilds, for example:
2BCl3 + 6LiH --> B2H6 + 6LiCl
The reaction would be performed by adding BCl3 to LiH suspended in an unreactive solvent such as ether or hexane, and the B2H6 condensed with liquid
nitrogen, or absorbed by something else depending on the intended use.
Boron trichloride seems fairly reactive so I am pretty sure the reaction will work.
I know the LiH + BF3 will produce B2H6 however BF3 isn't exactly a easy and cheap chemical to produce or buy. BCl3 on the other hand can be made
by reacting borax and charcoal with chlorine at high temperatures, which is somewhat complicated but at least doesn't require concentrated HF.
I have not seen any mention of this reaction anywhere, but don't see any glaringly obvious reason why it should work. Thoughts?12AX7 - 11-10-2005 at 12:06
Hmm, possible, but the given reaction could also result in a hydrogen reduction to boron in non-stoichiometric quantities, i.e., 2BCl3 + 6LiH = 2B +
6LiCl + 3H2.
Looks like it depends whether B + H2 or B2H6 is preferred.
Hrm.
That should be B2H4, shouldn't it Chris?
TimBromicAcid - 11-10-2005 at 12:41
B<sub>2</sub>H<sub>4</sub>...? Maybe you should brush up on your boron chemistry,
B<sub>2</sub>H<sub>6</sub> is quite the interesting moleulcule, suprised that in your knowledge database you haven't
taken a moment to examine its interesting structure and bonding.
Regarding the topic on hand it would be good to look at the energy of the reaction to see if it is comparable to the reaction using lithium hydride
and boron triflouride. I'm using values from the 1957-58 Handbook of Chemistry and Physics so they are different from the normal book values by
a constant.
[Note, I have seen the reaction as you describe it with lithium hydride and trifluoroborane but I have also seen it with sodium hydride, specifically
in Descriptive Inorganic Chemistry 3rd edition Rayner-Canham and Overton]
Really what you are comparing is:
NaF = -136.0
NaCl = -98.23
BF3 = -265.4
BCl3 = -100.0
Should work based on that but I can't do a delta G calculation because I can't find the literature values for some of the involved
molecules. However you have to consider what 12AX7 said about the possibility of it reducing it to elemental boron instead of the borane.
BTW, since I didn't mention it in this thread, diborane can be prepared by adding iodine dissolved in a glyme to sodium borohydride or similar to
give diborane in a good yield.garage chemist - 11-10-2005 at 13:26
I made a diborane- containing gas (contained lots of hydrogen, too) by adding 85% H3PO4 to sodium borohydride. It burned with a marvellous, brilliant
green flame and emitted a faint smoke of B2O3.
A method which has much better yields is to add an anhydrous solution of hydrogen chloride in ether to sodium borohydride.
The reagents must be anhydrous. Any water hydrolyzes the formed B2H6.
The liquid and solid borohydrides can be made by heating the resulting diborane which causes it to decompose into higher boranes and hydrogen. At
different temperatures, different boranes are produced. Pentaborane B5H9 is produced at 240°C according to Hollemann- Wiberg.
NaBH4 must be bought.
For people in Europe: Kremer now carries it. Enough said.
People outside Europe need to look for scientific suppliers.
[Edited on 11-10-2005 by garage chemist]BromicAcid - 11-10-2005 at 13:32
Quote:
NaBH4 must be bought.
There is a method of preparation that I have been considering, actually it is from a
patent so it likely isn't that reliable, I have terrible luck with patents afterall. Never the less, magnesium boride can be made by reacting
boric oxide with excess magnesium, not in great yields but >50% in good conditions. The patent relates that refluxing this magnesium boride with
alcoholic NaOH or KOH can give the respective borohydride in good yield and this can be in turn recrystalized from alcohol or even water (considering
the relative stabilities of the borohydrides to the short term action of water) I made a little note of this in the other thread on borohydrides I
believe and I plan to give it a try some day as it sounds like fun. But still it probably is a lot easier to just buy it.Eclectic - 11-10-2005 at 16:05
I made triethyl borane in college using a grignard reaction between EtMgBr and BCl3 in ether. Very exothermic, very flammable. The product is
spontaniously flammable and is the "green dragon" starter fuel for the SR-71 Blackbird.
Warning: Boranes are acutely toxic and spontaniously flammable.12AX7 - 11-10-2005 at 20:15
Huh, interesting, so it forms nonlinear groups, key to stability I'm guessing, unlike carbon which is stable when linear. So it's a dimer,
but not in the normal sense of the word, more along the lines of an amine salt. (I know, the electrons aren't arranged the same so you'd
say it isn't even close, but tetrahedral NH4+ is rather nonstoichiometric (4 bonds) as the nitrogen goes, too.)
Timgarage chemist - 12-10-2005 at 04:28
Does Triethyl borane spontaneously catch fire in air? Would be neat if it did.
I wonder if BCl3 can be made the same way as BF3, e.g. heat a mix of B2O3, CaF2 and lots of H2SO4 (this works, Hollemann- Wiberg lists it as a method
of preparation!).
One could substitute the CaF2 for NaCl and hope that it makes some BCl3 on heating.garage chemist - 12-10-2005 at 05:30
Wow! I found a route to BCl3 which is very feasible to do on a small scale and with easy apparatus.
Trimethyl borate is easily prepared from boric acid and methanol, with H2SO4 as catalyst.
Obvious problem here is the production of extremely toxic phosgene.
However, Ullmann offers a solution to this:
"A modification of this process, in which a free- radical initiator is used, has recently been reported (some sources are given).
A low reaction temperature, 40- 90°C, and the use of an initiator decrease the consumption of chlorine and prevent the production of
phosgene."
So we just need to find a suitable initiator. A peroxide compound should do the trick.
[Edited on 12-10-2005 by garage chemist]Chris The Great - 12-10-2005 at 11:09
Very cool, much easier to do then heating at 700 degrees!
And it gives phosgene which can be useful as well, as long as you actually want it! I certainly wouldn't want to make it accidently!
Does it say what the reaction would be if catalysed properly, or do have any idea? I'm not really an expert on the chemistry of boron, I'm
just learning at the moment.garage chemist - 12-10-2005 at 12:14
You want to avoid phosgene formation completely, because BCl3 and phosgene cannot be separated by distillation (boiling points too close together).
Maybe by forming the high- boiling complex with diethyl ether the BCl3 can be separated from the phosgene?Chris The Great - 12-10-2005 at 23:11
Yeah I realized that a little while later, to my dissapointment. I can't find any solubility information of BCl3 either, so I don't know
wether the gas could be bubbled through a suitable solvent to pick up one of them and let the other continue through. The complexing idea would seem
to be the best if they where to be seperated.
I was reading PATR and it mentioned that diborane can explode when heated above it's bp (I would assume under pressure). However I have never
seen any mention to problems with handling liquid diborane during rocket research and it is an important component for manufacturing silicon computer
chips, so has anyone else heard this from another source? It seems unlikely as diborane is only a very small positive heat of formation and so an
explosion would not be that energetically favourable. Perhaps hydrogen release as large borane compounds are formed?Nerro - 23-5-2006 at 11:39
BCl3 should be quite soluble in benzene AFAIK.
And no it does indeed not usually form linear molecules due to the electron deficient nature of B. It forms clusters based on wades rules. which
basically states that boranes form "aromatic" cages with BH clusters on the vertices. The cages are formed when the general formula is
B<sub>n</sub>H<sub>n+2</sub>. These clusters are called "closo" molecules. There are also "nido" molecules with the general
formula B<sub>n</sub>H<sub>n+4</sub>; "Arachno" structures with the general formula
B<sub>n</sub>H<sub>n+6</sub> and "Hypho" molecules with the general structure
B<sub>n</sub>H<sub>n+8</sub> although the latter are not very common or stable. Stability decreases from closo to hypho with
closo being quite stable and hypho being quite unstable. In the closo molecules the "+2" Hydrogens are never present which is why the closo molecules
are allways divalent anions.
for more information search for "Wade's Rules" or "Boron chemistry" or "closo +nido + arachno + hypho" Chris The Great - 23-5-2006 at 18:16
I saw something today, called diboramine. The formula was B2H7N. This was in Chemistry and Technology of explosives. Apparently the bp is around 70
degrees or so. I thought it might be a somplex between diborane and ammonia, but that doesn't add up. I've also not seen anything from googling,
expect borane amine is BNH6 and a solid at STP, nothing about "diboramine".
It was in a section on rocket fuels.kazaa81 - 29-5-2006 at 05:37
Quote:
Originally posted by garage chemist
Does Triethyl borane spontaneously catch fire in air? Would be neat if it did.
How can triethyl borane be made? By dissolving boric acid (H3BO3) in ethanol (C2H5OH) and the adding a lewis acid like sulphuric acid to catalyze the
reaction?12AX7 - 29-5-2006 at 06:55
That would make an ester, not a borane. You'd need like, a grignard or something.
Timneutrino - 29-5-2006 at 08:22
>By dissolving boric acid (H3BO3) in ethanol (C2H5OH)...
No, that will give you triethyl bor<b>ate</b>:
B(OCH<sub>2</sub>CH<sub>3</sub><sub>3</sub>
not triethyl bor<b>ane</b>:
B(CH<sub>2</sub>CH<sub>3</sub><sub>3</sub>garage chemist - 29-5-2006 at 10:00
Please, read the thread! The method for making triethyl borane has been covered here. Just another case of not reading before asking.kazaa81 - 30-5-2006 at 03:57
Here are some interesting information I've found about triethyl borate...from CRC Handbook of Chemistry and Physics 85th ed.
Triethyl Borate [boric acid, triethyl ester]
Chem. Formula: B(OCH2CH3)3 or C6H15BO3
CAS #: 150-46-9
Mol. 145.992 g/mol
Form: liquid
M.P. -84.8°C
B.P. 120°C
Denstity 0.8546^20 g/cm3
Miscible in ethanol
So, would I be able to separate it from ethanol, by boiling it off and condensing? --> yes
Back to topic please!!! Boric acid esters have nothing to do with boron hydrides!
[Edited on 30-5-2006 by chemoleo]gthulasi - 10-7-2006 at 16:27
Can anyone say the decomposing temperature of Boron Haldies BF3,BCl3
