Sciencemadness Discussion Board - Isolation of Boron from Borax (Sodium Tetraborate)

Isolation of Boron from Borax (Sodium Tetraborate)

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bismuthate - 28-9-2013 at 13:46

would a solution of borax + chlorine make boron oxide because that would be an easy route to elemental boron

[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: replaced subject "boron" to reflect topic of merged threads]

[Edited on 14.10.13 by bfesser]

12AX7 - 29-9-2013 at 03:02

Use HCl, not chlorine.

B(OH)3 will precipitate. It has to be fused to make boron oxide.

Tim

bismuthate - 29-9-2013 at 03:23

i'm not allowed to use heat to decompose it so i was wondering can it react with chlorine?

blogfast25 - 29-9-2013 at 05:57

Quote: Originally posted by bismuthate

i'm not allowed to use heat to decompose it so i was wondering can it react with chlorine?

Reacting it with chlorine won't do zilch.

Borax is a (complex) salt of a very weak acid, i.e. boric acid (B(OH)3). When a stronger acid (like HCl) is added to a solution of borax the stronger acid displaces the weaker one. Simply put:

Borax (solution) + HCl (solution) === > B(OH)3 (oily flakes) + NaCl (solution)

There are plenty recipes for this, use the search facility.

Then filter off the boric acid, wash and dry it. Then, on strong heating, it loses water and forms a glassy mass of B2O3:

2 B(OH)₃(s) ==(heat)== > B₂O₃(s) + 3 H₂O(g)

This can be reacted with Al or Mg powder to produce elemental boron. But obtaining pure boron this way isn't easy. See Brauer (library) for a procedure.

If you're not allowed to use heat, then forget about chemistry: heat is what drives much of it!

[Edited on 29-9-2013 by blogfast25]

bismuthate - 29-9-2013 at 06:42

well i can use heat for short times via alcohol burner and reactions but i can't use a large fire or a stove which is used to decompose boric acid. is there any other way to make it like a way to make BCl3 or something along the lines of that? (probably not i can't think of a way)

blogfast25 - 29-9-2013 at 07:48

An alcohol burner might just be hot enough to slowly dehydrate boric acid. But I'm not sure what temperature is needed for FULL dehydration (obtaining anhydrous B2O3) Smart use of your burner and correct choice of crucible (copper for instance) could do wonders here.

BCl3 from borax is a hat trick and you'd need to go via B2O3 anyway. It's extremely reactive towards water...

[Edited on 29-9-2013 by blogfast25]

bismuthate - 29-9-2013 at 07:55

i have a idea! i can put the crucible of borax on a Al/MnO2 thermite mixture which will run for quite a while. will that work?

blogfast25 - 29-9-2013 at 08:46

Quote: Originally posted by bismuthate

i have a idea! i can put the crucible of borax on a Al/MnO2 thermite mixture which will run for quite a while. will that work?

NO! Thermite mixtures, ESPECIALLY MnO2 thermites, burn far too hot and far too short.

Try this. Use your spirit burner on a 'high' setting (longish wick), with methanol if possible.

Now you'll need some kind of a 'crucible'. A piece of copper sheet would be ideal but if you haven't got any some mild steel sheet cut from a steel pop can for instance, could work. Fashion it into a concave surface, some 3 x 3 cm maximum. Load a quarter of a teaspoon of the dried boric acid onto it. Now try and somehow clamp this 'crucible' at the right height above the burner, so you don't have to hold it (which would quickly become tiresome) and light up the burner. Or try a disused teaspoon (thin metal if possible) as a crucible...

Carefully observe the boric acid: it will probably melt and should start giving off steam. Whether you'll reach high enough temperature to fully dehydrate it to B2O3 I don't know. It will in any case take time, at least 30 min on full heat.

On p 57 of 'The Golden Book of Chemistry', which you can download here:

http://chemistry.about.com/b/2012/10/01/download-the-golden-...

... is described an experiment that claims to prepare boric oxide from boric acid, with a teaspoon and a spirit burner. But it's not clear to me whether actual B2O3 or rather one of several intermediates (HBO2 or H2B4O7) is obtained. It's largely a question of heating time, I'm guessing.

[Edited on 29-9-2013 by blogfast25]

bismuthate - 29-9-2013 at 12:34

That would work, but could I silver instead of copper for better conductivity?

elementcollector1 - 29-9-2013 at 12:41

Quote: Originally posted by bismuthate

That would work, but could I silver instead of copper for better conductivity?

Why would you want better conductivity?
If you mean heat conductivity, I think you'd want less, not more: More heat conductivity will mean more heat is dispersed to the atmosphere.

papaya - 29-9-2013 at 12:48

Look at this
https://www.youtube.com/watch?v=3IAv6Ohjuh8
It may be that one of the products is boron (I'm not sure).

bismuthate - 29-9-2013 at 14:43

What was that? (I only speak English, sorry)
Also, could i burn a pile of Mg ribon to provide heat?

violet sin - 29-9-2013 at 15:22

I don't know that every one has these available, but the little stainless steel condiment cups are great for stuff like this. I went to a near by wall-mart and got 2 four packs of em for 97cents each. practically disposable at that cost.

http://www.walmart.com/ip/Mainstays-Condiment-Cups/14964954

work great for me

bismuthate - 29-9-2013 at 15:32

Thanks for the tip. Also, kind of off topic but i went to Stop and Shop and i saw 100ml pyrex custard cups and they would make great beakers for small experiments.

bfesser - 29-9-2013 at 16:22

Beware that new "Pyrex" cookware isn't borosilicate, and therefore prone to thermal shock induced shattering.

Tdep - 29-9-2013 at 16:35

Is there any distinguishing way to tell if glassware is borosilicate or not? Does it have to be labelled as such or is your only clue when it shatters under heat shock?

bismuthate - 29-9-2013 at 16:36

i wasn't suggesting usin pyrex for this reaction, i was just mentioning somthing i saw at the store. sorry for the confusion.

bfesser - 29-9-2013 at 18:15

The cookware issue was discussed a few years ago in <a href="viewthread.php?tid=12170">several topics</a>.

[Edited on 30.9.13 by bfesser]

papaya - 29-9-2013 at 20:26

Quote: Originally posted by bismuthate

What was that? (I only speak English, sorry)
Also, could i burn a pile of Mg ribon to provide heat?

It is a mixture of borax + magnalum, a non-standard thermite, I suppose the boron (or some borides) may be on of the products (another is sodium as you can see it afterburning in video).

bismuthate - 30-9-2013 at 03:37

thats interesting if i mix borides with HCl i get boran which i can burn to get B2O3 right?

blogfast25 - 30-9-2013 at 04:53

Quote: Originally posted by elementcollector1

Why would you want better conductivity?
If you mean heat conductivity, I think you'd want less, not more: More heat conductivity will mean more heat is dispersed to the atmosphere.

No. The highest temperature of a material in a rudimentary crucible (all other things being equal) is obtained with the highest heat conductivity crucible material and the thinnest crucible wall thickness. That's basic heat transfer science. There's no getting around that.

[Edited on 30-9-2013 by blogfast25]

papaya - 30-9-2013 at 06:56

Quote: Originally posted by bismuthate

thats interesting if i mix borides with HCl i get boran which i can burn to get B2O3 right?

In theory yes I think - see videos on the same channel - with phosphate instead of borax he gets something that when water(or acid) is added releases a gas immediately catching fire - phosphine. (oh - easy way to P2O5

)

bismuthate - 30-9-2013 at 11:20

that sounds awesome shame i don't have any phosphates

would this work with silicates? just out of curiosity.

papaya - 30-9-2013 at 11:31

search for SiO2 + Al + S thermite. Btw. there's an entire discipline in material synthesis (mostly ceramics) called SHS
https://en.wikipedia.org/wiki/Self-propagating_high-temperat...

bismuthate - 30-9-2013 at 11:38

actualy i was talking about silicates mixed with aluminum. (I have already seen many videos about sand thermite) hovever i found that link very interesting. maybe i'll use one of those reactions soon.

[Edited on 30-9-2013 by bismuthate]

papaya - 30-9-2013 at 11:50

Well, you could try to do Mg + Na2SiO3 reaction and report back the results if you want the silicate (what about glass+Mg?). I don't know the outcome of course, don't forget that all these boranes, phosphines, silanes are a fire hazard and very toxic.

bismuthate - 30-9-2013 at 12:04

i would love to i have some pyrex an mg ribbon that should work
maybe we coulb make boon from glass!

blogfast25 - 30-9-2013 at 12:22

Quote: Originally posted by papaya

)

I hope that latter bit was an attempt at irony: phosphine and diphosphine are very, very toxic, as well as extremely flammable. From Wiki:

NIOSH recommends that the short term respiratory exposure to phosphine gas should not exceed 1 ppm. The Immediately Dangerous to Life or Health level is 50 ppm.

Trying to prepare B2O3 from burning boranes is folly when you can just heat boric acid to full dehydration. Borides aren't that easy to prepare either.

blogfast25 - 30-9-2013 at 12:28

Quote: Originally posted by papaya

Well, you could try to do Mg + Na2SiO3 reaction and report back the results if you want the silicate (what about glass+Mg?).

Huh??? Fine sand and an excess Mg powder can be reacted to Mg₂Si, magnesium silicide, which when treated with acids like HCl gives a mixture of silanes, mainly monosilane and disilane. There must be numerous Utoobs on Mg₂Si and silanes. They do spontaneously combust when in contact with air...

[Edited on 30-9-2013 by blogfast25]

bismuthate - 30-9-2013 at 12:28

i want to see if i can make boron from glass as an unusual method.

blogfast25 - 30-9-2013 at 12:32

Quote: Originally posted by bismuthate

i want to see if i can make boron from glass as an unusual method.

Even borosilicate glass only contains about 15 % B2O3. And glass is a notoriously unreactive substance. Honestly, try and make some B2O3 in the usual way and look up in Brauer's 'Preparative Chemistry' (see library) how to get boron from that.

[Edited on 30-9-2013 by blogfast25]

bismuthate - 30-9-2013 at 12:38

i know how you mix B2O3 with mg and ignite it then wash with HCl, but i want to prepare it with a unique method.
lithium reacts with glass and i have some.

Reaction of borax decahydrate with an acid

Upsilon - 12-10-2013 at 19:22

I am attempting to isolate a sample of boron from borax, and every source I find says that the first step is to react it with some acid. Unfortunately, I am having extreme difficulty coming up with a balanced equation for the reaction of borax decahydrate with an acid. I would like an equation for this reaction using HCl and for using H2SO4, just for reference. I know the decahydrate portion plays a significant role in the reaction, but I just can't figure it out.

mayko - 12-10-2013 at 19:59

This should help

http://en.wikipedia.org/wiki/Boric_acid#Preparation

WGTR - 12-10-2013 at 20:37

Just a suggestion...boric acid can usually be purchased at the grocery store as "roach killer". It's pretty cheap.

violet sin - 12-10-2013 at 20:54

It's fun to mix the boric acid w/ methanol to make green flames

trimethyl borate. It's not a crazy reaction that goes off by it self, ya have to light it. I have had some in a bottle since Halloween 2 years back. I used the roach powder from Walmart. Cheap fun, but it does burn fairly hot so watch out. I had put some in a beer bottle and once the flame got me. Nothing bad, just missing hair on my hand

I know that this wasn't the question you were asking but with Halloween coming up, why not mention it. For the methanol I used "HEET", yellow bottle found in auto isle for flushing radiator systems. It's also fairly cheap at Walmart, ~2$ each. I got a couple 4 packs ther and am still using it for various things.

Woulda put in the links to specific items I describe, but I'm outta town working again and only have an iTouch for Internet connectivity, sry.

[Edited on 13-10-2013 by violet sin]

Upsilon - 12-10-2013 at 20:58

Ah, thanks. I have seen that equation before but I was trying to use an equation balancing applet to confirm it, but I guess it doesn't deal well with hydrates. Also, is the roach killer pure boroc acid? It seems a little unlikely. Now, onto isolating some boron...

I'm trying to see if it is possible with cheap, common household ingredients. First would be to react borax with excess vinegar at elevated temperatures with this equation:

Na2B4O7 • 10H2O + 2C2H4O2 -> 2Na2C2H3O2 + 4H3BO3 + 5H2O

Then, the resultant boric acid is heated to decompose into boric oxide. The boric oxide is then made into a solution and aluminum foil is added:

B2O3 +2Al -> Al2O3 + 2B

I have checked and aluminum is capable of replacing boron in single replacement in an aqueous solution at room temperature.

Now, this procedure has some difficulties. After reacting the borax with vinegar, there will probably be a big pile of a mixture between leftover borax, sodium acetate, and boric acid. I read that if the solution is cooled back down, the boric acid will precipitate leaving the other salts dissolved in water, although I'm not sure of the validity of this statement. Also, the end product will be a solution containing boron and aluminum oxide; how could I separate them?

violet sin - 12-10-2013 at 21:12

I have found pure boric acid in CVS and Rightaid stores. Antiseptic wash aid I believe, came as a dry powder with USP(like other things used on people). The roach powder was blue tinged and not pure. But Much cheaper( not the point I understand). The stuff @ CVS was like 7$ for a medium sized pill bottle worth.

For your plans borax may be fine. Could you buy a small bottle of HCl? 3$ at ACE hardware. I seem to remember seeing some vids on YouTube a while back. Borax was more soluble than boric acis, so when HCl was mixed in the BA precipitated out. Obviously some was still in sol. But who cares.

There were also vids there of people using the thermite rxn with Al powder under a coffee can. They had heated their boric acid to get the oxide, crushed it and mixed in Al. After it was done the guy used acid(again HCl I think) to etch off the crud. And was left with black shiny lumps. I'll see if I can come up with a couple vids for ya, but take a look your self too.

Hope that was a bit more productive of a response
http://m.youtube.com/watch?v=0QBCyOrjR2o
http://m.youtube.com/watch?v=0kbm92a2ltc

[Edited on 13-10-2013 by violet sin]

[Edited on 13-10-2013 by violet sin]

WGTR - 12-10-2013 at 23:48

The roach powder that I found was 99% pure. I recrystallized it from hot water a couple of times because I was using it in electroplating (purity is important in that case).

I have a thought question for you. What is being driven off from the boric acid in order to make it into boric oxide?

Also, if you weaken aluminum's tenacious oxide layer (with gallium or mercury), does it react with water?

weiming1998 - 13-10-2013 at 01:47

When you heat boric acid to make boron trioxide, water vapour is being driven off.

And aluminium indeed reacts with water when its oxide layer is removed with gallium or mercury. This should demonstrate it pretty well.
https://www.youtube.com/watch?v=JasZ8V6LpbQ

Upsilon - 13-10-2013 at 06:23

Even though pure aluminum will react with water, I still have seen demonstrations of putting aluminum foil in a CuCl2 aqueous solution, yielding copper metal (even though some hydrogen gas was still produced).

Also, how concentrated is the HCl from the hardware store? I know I can also get muriatic acid (HCl) from the local pool store. Would vinegar not just do the same job anyway, just slower? Even if it is only a 5% solution, it should still work if I add enough. I guess I'll try this out and share my results.

[Edited on 13-10-2013 by Upsilon]

violet sin - 13-10-2013 at 07:46

Muriatic is fine, that's actually what I meant. The acid has to be strong enough to take the sodium and protonate the weaker acid from borax. Also at only 5% sol you are giving it too much water to dissolve your boric acid product I think. Prob not goin to hit its sol limit and ppt.
Hope it helps, but I literally just woke up so hope I explained it right

bismuthate - 13-10-2013 at 07:47

I would not recomend vinegar because it is 95% water which the boric acid can dissolve in, effectively lowering your yield. I belive that concentrated acids are best in this situation. However if your yeild doesn't matter it would work.

WGTR - 13-10-2013 at 11:32

Quote: Originally posted by Upsilon

Even though pure aluminum will react with water, I still have seen demonstrations of putting aluminum foil in a CuCl2 aqueous solution, yielding copper metal (even though some hydrogen gas was still produced).

If boric oxide is simply dehydrated boric acid, then when you add boric oxide back to the water again, does it reform boric acid? And what would happen if you add aluminum to that?

The videos that I have seen so far involve a thermite-type reaction under anhydrous conditions. It would be interesting to see if you can do it in an aqueous solution.

MichiganMadScientist - 13-10-2013 at 12:45

Not sure if this helps, but Rite Aid pharmacies selling over the counter pure boric acid for like $5.00. It was located near the iodine/antiseptics section of the store.

http://shop.riteaid.com/Rite-Aid-Boric-Acid-Powder/dp/B005X7...

Upsilon - 13-10-2013 at 13:47

There are no Rite Aid stores near me, and I can't seem to find any boric acid on the CVS website. Anyway, as for my experiment so far, I added 150 mL of vinegar to 10g of borax, heated it, and let it sit for a while. I then refrigerated it hoping to see some precipitate, but there was none. Now I just set the solution outside to let the vinegar evaporate to leave a mixture of borax hopefully along with some sodium acetate and boric acid. I'm just going to dump the mixture in some water, dissolve it, toss in some aluminum, and see if any black crystals form.

bismuthate - 13-10-2013 at 14:09

There was no percipitate because excess water dissolved the boric acid.
Also aluminuium will not percpitate boron first because the aluminum pssivates and second because if they did react alumminium borate would be formed, not boron.

Upsilon - 13-10-2013 at 15:39

Quote: Originally posted by bismuthate

There was no percipitate because excess water dissolved the boric acid.
Also aluminuium will not percpitate boron first because the aluminum pssivates and second because if they did react alumminium borate would be formed, not boron.

Aluminum certainly does have the potential to replace boron, I'll pull up my source if you need me to. Also, would you please show the balanced equation for the formation of aluminum borate? Some may form, but there would still be an ample amount of boron. People usually extract boron from boric oxide by reducing it with magnesium, but aluminum would work just as well.

bismuthate - 13-10-2013 at 16:28

I did not say it did form I said that it would form if there was any reaction. However I very well may be wrong. Please do show this reference.
Also yes it is reduced with magnesium but from my experience never while in solution.
ps The equation would be 2Al+2H3BO3==>2AlBO3+3H2 however this reaction does not take place due to passivisation.

Upsilon - 13-10-2013 at 16:51

I do not have the exact website, however it is a PDF file called "redpot.pdf" at courses.chem.indiana.edu

Edit: Ah, I found it. Here it is: https://www.google.com/url?sa=t&source=web&cd=1&...

Sorry about the huge link, but I'm on a mobile device and I cannot get the exact link, I just have to use the Google search link.

[Edited on 14-10-2013 by Upsilon]

bismuthate - 13-10-2013 at 17:13

Thanks. Also I have few thoughts.
1st Could a Ga Al alloy allow Al to react with boric acid?
2nd I am sorry if the equation was incorrect there seem to be many types of aluminium borate.
3rd I know that using a mobile device is a pain I typed this with a kindle with no keypad. It took forever.

[Edited on 14-10-2013 by bismuthate]

Upsilon - 13-10-2013 at 18:13

Ah, I see where we had a miscommunication here. You very well could be right, but I was not referring to reacting boric acid with aluminum. I intended to heat the boric acid to decompose it into water vapor and boric oxide, THEN reacting it in a solution with aluminum. Sorry if that wasn't clear.

WGTR - 13-10-2013 at 21:07

You might like to consider, that if boric acid is heated to dehydrate it to boron trioxide; then by putting it back into a water solution, you are rehydrating boron trioxide to boric acid again. Boron trioxide doesn't exist in an aqueous solution. The Wiki page for boron trioxide is a bit deceptive, because it gives a solubility figure for it. In actuality, it is reacting with water to form boric acid.

If you look at the solubilities listed for boric acid and boron trioxide, the solubilities for boron are almost the same in both cases.

B2O3 solubility: 22 g/L (at unknown temperature).

(1 mol B2O3/69.6182 g B2O3) x (22 g B2O3/1 L) x (2 mol Boron/1 mol B2O3) = 0.63 mol Boron/L

H3BO3 solubility: 4.72 g/100 mL (at 20°C).

(1 mol H3BO3/61.83 g H3BO3) x (4.72 g H3BO3/100 mL) x (100 mL/ 0.1 L) x (1 mol Boron/1 mol H3BO3) = 0.763 mol Boron/L

If the B2O3 solubility was measured at a lower temperature than 20°C, then that could account for the small difference in calculated solubilities.

Anyway, those simple algebraic equations shown above were almost the first things learned in college chemistry. If you read up on Stoichiometry, you'll get an idea of how to set up these types of equations.

[Edited on 14-10-2013 by WGTR]

bismuthate - 14-10-2013 at 03:10

Sorry for the confusion before. Maybe if you wanted to do the experiment you could find a solvent B2O3 is soluble in or if there is none dissolve the aluminum in mercury. (quite an odd idea). The main rule is you need to keep the water away from your B2O3 because that will couneract all the time spent dehydrating your boric acid.

Upsilon - 16-10-2013 at 09:50

I guess I see why the aqueous single replacement won't work; I had been wondering that myself. If I reduce the boric oxide with no solvent, then the boron sample would be significantly contaminated. I assume it is possible to dissolve boric acid in propylene carbonate? This would make it easier to obtain a pure sample.

bismuthate - 16-10-2013 at 13:21

In the begining of this thread I was told that borax does not react with Cl. This raise the question "how does it react with fluorine"? Under anhydrous conditions of course.

Upsilon - 16-10-2013 at 13:45

Quote: Originally posted by bismuthate

In the begining of this thread I was told that borax does not react with Cl. This raise the question "how does it react with fluorine"? Under anhydrous conditions of course.

I'm going to guess there is some reaction with fluorine, but it depends on the reduction potential of tetraborate. For some reason it is exceedingly difficult to find information on the reduction potential of polyatomic ions; in fact, I have yet to get any information on the subject. Regardless it is much more practical to produce boric acid from borax + an acid.

bismuthate - 16-10-2013 at 14:02

This was not meant to be practical it was only scientific curiosity. I would not want to deal with fluorine.

Upsilon - 16-10-2013 at 18:53

Well anyway like I said you'd have to find a source telling you about the reduction potentials of polyatomic ions, and I wish you luck in your search.

AJKOER - 18-10-2013 at 17:18

An extract from Atomistry.com (link: http://boron.atomistry.com/ ) on Boron and its preparation:

"Amorphous boron is obtained by passing the vapours of the chlorine compound over heated sodium, or, quite similarly to silicon, by igniting the oxygen compound with magnesium. After the removal of the admixtures, it forms a black powder of the density 2.5, which in many respects behaves similarly to charcoal, but is more easily oxidised; this occurs more especially by means of strongly oxidising solutions even at the room temperature.

By the fusion of boron trioxide (vide infra) with aluminium, crystallised boron is obtained, which, on account of its hardness, has been called " adamantine boron." It is not obtained quite pure in this way, but contains aluminium derived from its preparation. Since this metal is the element most nearly related to boron, the product is not to be looked upon as a compound, but as a mixture (possibly with a diamond-like form of aluminium isomorphous with boron, and not known by itself).

Boron containing carbon, and obtained from the two elements at a very high temperature, is of a similar character, and also possesses an adamantine hardness. This also ought most probably to be regarded as a solid solution, and not as a chemical compound."

bismuthate - 19-10-2013 at 04:54

Thanks! The boron chlorides method gave me an idea. Could trimethlyborate (boron timethoxide) be reduced by sodium (or lithium) to yield sodium methoxie and boron or am I missing the obvious?

[Edited on 19-10-2013 by bismuthate]

Upsilon - 19-10-2013 at 16:35

Just out of curiosity, what does borax decompose into when electrolyzed in an aqueous solution?

bismuthate - 19-10-2013 at 16:47

In solution NaOH and H3BO3 I would believe. However the two will quikly react.
P.S. it isn't decomposition.

Upsilon - 19-10-2013 at 17:15

If you were to use a Hoffman apparatus, however, wouldn't you be able to keep the solutions of NaOH and H3BO3 separate?

bismuthate - 19-10-2013 at 17:25

Yes, but it is more expensive than using HCl. Cool way to make NaOH though.

Upsilon - 19-10-2013 at 17:32

I already have a Hoffman apparatus I made out of PVC. Are you saying that the electricity will be more expensive than the HCl?

bismuthate - 20-10-2013 at 02:55

I'm not sure about you but I have a liter of acid for $2.

WGTR - 20-10-2013 at 08:32

You can isolate boric acid by electrolysis, but you need a suitable diaphragm to keep the anode and cathode compartments separate, and you need a suitable non-soluble anode.

If you use a cathode compartment that is much larger than the anode compartment, this will make it easier to select a diaphragm. This is because it limits the concentration of the NaOH in the cathode compartment to a low level. Industrially, obtaining concentrated NaOH with a diaphragm cell is difficult, because the caustic effects of the concentrated hydroxide attack many materials that would otherwise be good diaphragms. At the same time, borax is not very soluble, so your final concentration of NaOH probably wouldn't be too high even if the two compartments had equal volume. Real cellophane might be a suitable cheap option for you.

If you start with a saturated solution of borax, as electrolysis proceeds your precipitate in the anode compartment should consist more and more of boric acid.

The anode will have to be something that will not react under the cell conditions. Platinum would work great. You could experiment with other materials to see what works for you.

Reacting HCl with borax would probably be much easier (and faster) than separating it out by electrolysis. Or just buying the boric acid from the grocery store.

Upsilon - 20-10-2013 at 10:52

For my electrolysis experiments, I use graphite rods that I extract from common pencils. They hold up quite well, and it is no big deal of they need replaced since they're so inexpensive.

Upsilon - 20-10-2013 at 13:44

Well I went to Lowes and picked up 2 gallons of 31.45% HCl for around $10. Will be experimenting soon.

avrpunk - 20-10-2013 at 14:52

Upsilon, graphite from pencils is quite expensive compared to a welding supply store which will have "gouging rods" in varying diameters. The only preparation needed is to remove the copper coating from the part that will be immersed, which is quite easy given that you're doing electrolysis in the first place, and the rest is graphite.

Upsilon - 23-10-2013 at 11:04

Alright, so I started with 100g of borax and added excess HCl and stirred. Afterward, after about 5 minutes the solution settled into a prominent green-yellow layer floating above an opaque, thick white layer (about 3 times larger than the green layer) that I'm assuming is an extremely saturated NaCl solution with excess NaCl floating around. I'm going to let it sit for a while yet longer and see what happens. Then I'll remove the HCl layer and heavily dilute the other layer to hopefully yield some amount of boric acid.

UPDATE: After removing the excess HCl, I added more borax to get rid of any residual HCl, periodically throwing in a pinch of sodium carbonate to test the presence of HCl. I also added some nom-distilled water to dilute the solution in hopes to get a better view of my product. What I see now is a large top layer of frothy white solid above a layer of water. Underneath the water is another layer of white solid. Neither boric acid nor NaCl is less dense than water, so what is this solid that is floating to the top?

[Edited on 23-10-2013 by Upsilon]

violet sin - 23-10-2013 at 11:53

Did you check solubility data on NaCl and B(OH)3? I don't think the thick opaque stuff is NaCl. You should let the ppt. settle out, decant off HCl/NaCl liquid and wash the ppt. W/ ice cold Distilled water.

Diluting it will dissolve more of your product into the waste solution.

NaCl sol. data from wiki...... 359g/100ml ( no temp given )
B(OH)3 sol..."...."....."......... 2.52g/100ml @ 0'C
........................................ 27.53g/100ml @ 100'C
Borax sol....."....."....."....... No numbers give but stated easily sol. in water

The equations given in wiki for conversion products is 4mol B(OH)3 and 2 mol NaCl. It stands to reason 2 mol of highly soluble NaCl will stay in sol. But the 4 moles of low sol. boric acid won't. ie your white ppt is boric acid no?
*************************************
But CO2 bubbles on small boric acid crystals is lighter than water, probably your floaters. If you started with just enough water to dissolve the borax( add stirr, repeat. not super critical ) then calculated your HCl to be a slight excess needed to convert the borax. Add and wait for rxn. Calculate NaCl moles produced, and reduce volume of sol. by heat till the salt is nearly saturated. And then chill to ppt. max B(OH)3, filter off, wash w/ ice cold water, you should have fairly clean boric acid. You could then dissolve it in just enough boiling water and wrap you beaker in a towel for insulation so cooling is slow. And get larger, easier to filter/wash nice boric acid.
[Edited on 23-10-2013 by violet sin]

[Edited on 23-10-2013 by violet sin]

Upsilon - 23-10-2013 at 12:18

I skimmed the surface of the settled solution with a spoon and gathered the floating substance. I put it in a new beaker and added a few hundred ml of distilled water. Now only a fraction of the solid is floating above the water while the majority settled at the bottom. I'm going to repeat this process and see if any still floats.

violet sin - 23-10-2013 at 13:06

You will be able to get boric acid the way you did, but I think it would be more rewarding to crunch the numbers and all next time. I did a quick run through and if you started with 100g borax fully dissolved you would aprox 52ml of the 31.45% HCl, + throw in a slight excess. After rxn that would make aprox 30.34g NaCl, 61.83g B(OH)3. You could calculate aprox yield off theoretical in the end. No guessing on neutralization HCl, baking soda not needed. Heat to reduce sol. volume and bonus drive off excess HCl. Though it would be hard to reduce it till NaCl saturation as I had said, would be near nothing left liquid wise with the boric ppt at bottom and only 30.34g NaCl. So just reduce it enough to not spend all day gravity filtering for first isolation of boric acid. Then try the boiling water/towel on beaker thing( worked for me in the past quite well just haven't tried on boric acid) for larger crystals and less drain time. Wash with ice cold water a few times and should be nice and clean.

There might be a slight mistake or two on the exact numbers, I calculated the needed HCl from weight/weight % of 1000ml @ 1.16g/ml. Wasn't labeled strictly w/w % on MSDS but I don't think they would do w/v or volume/volume % with a gas dissolved in liquid. But I don't have a bottle handy to check( its back home with my pool).

Man it sucks typing this up on an iTouch ): hope it helps

Upsilon - 23-10-2013 at 13:23

I did the number crunching first but ended up just winging it during the experiment, as the calculated volume of HCl didn't seem quite enough. Anyway, I think it is safe to say that the substance I skimmed off of the surface is relatively pure boric acid. If I put a pinch of sodium carbonate in the solution, it floats to the bottom and when it reaches the solid bottom layer, it goes up in a mini white mushroom cloud, my guess is that some material is being carried up by the CO2 that is formed.

WGTR - 23-10-2013 at 13:34

The solubility of Na2B4O7 is comparable to H3BO3:

H3BO3 (g/100mL):

Due to the similar solubilities of these two compounds, in order to force the H3BO3 to precipitate without Na2B4O7, excess H+ should be added. This can be accomplished by adding excess HCl. However, if too much is added, then NaCl will also precipitate, due to the excess Cl-. NaCl is much more soluble than H3BO3 or Na2B4O7, however:

I would suggest starting with these quantities:
Na2B4O7: 0.04 moles = 8.0 g anhydrous, or 15 g decahydrate.
HCl: 1.2 mole = 120mL of 32% HCl.
Add all of this into water, forming 1 liter of solution at 20-25C.

Everything should dissolve.

Now cool the solution down to 0C on ice.

This should yield about 0.08 moles (4.9 g) of pretty pure H3BO3 precipitate, if I haven't missed something. The 0.04 moles of NaCl should stay in solution.

After your first batch, you can successively add extra Na2B4O7 to the warm solution to retrieve more H3BO3 on cooling. Eventually NaCl will build up in the solution, giving product that is not as pure.

I suggest making sure that everything dissolves on adding Na2B4O7 just to be sure that the system reaches equilibrium before you collect your precipitate.

[Edited on 23-10-2013 by WGTR]

violet sin - 23-10-2013 at 14:53

WGTR
Wiki:Na2B4O7·10H2O + 2 HCl → 4 B(OH)3 [or H3BO3] + 2 NaCl + 5 H2O

The suggested ratio was 1 x borax + 2 x HCl --> 4 x boric acid + 2 x salt( aka NaCl) so... Why are you suggesting 1.2 mol HCl for 0.04 mol borax? That is not a slight excess. Compared to borax, NaCl solubility and boric acid precipitation would drive rxn to the right any way, no? The weak acid sodium salt (borax)should readily hand over its sodium to the strong acid forming NaCl. unless the common ion solubility forces NaCl ppt like you stated it stays in sol. also 0.04 compared to 1.2mol is really not advisable right? For theoretical you would only need .08 mol HCl... Your suggested 120ml is more than twice what I recommended for one hundred grams of borax. Also there would be 0.08 mol NaCl not 0.04.

While reducing temp to ppt boric acid, NaCl isn't going to try and make a strong acid and borax again. Simple strong acid/ weak acid salt swap rxn. I don't think even by heating to reduce volume, that you could drive HCl back out forcing borax formation again.

Your info suggests heating initial water, total dissolution of borax, addition of SLIGHT excess HCl, cooling to force B(OH)3 ppt, and filtering. Not way over saturation of HCl.

At 0'C in 1 liter of water can retain ~25g boric acid. How ya going to get that to ppt out again? ( Sol data from wiki a few posts back boric acid 2.52g/100ml @ 0'C)

Just looked at your charts. How is >500g/l @ 100'C borax compare to 35g/l(?) @ 100'C for boric acid comparable? Side by side the curve is similar but the scale is not even close. Really not trying to be a dick here, so don't get offended. But looks like you may have made some inaccurate calculations or assumptions. Correct me if I'm wrong.
[Edited on 23-10-2013 by violet sin]

[Edited on 23-10-2013 by violet sin]

[Edited on 23-10-2013 by violet sin]

bismuthate - 23-10-2013 at 14:55

Idea: Make boron by passing BCl3 over lithium.
How: Take an excess of Boric acid and add it to methanol then distill the trimethyl borate at 70 degrees celcius. Boil the distillate and pass the vapor over lithium hydride. Then take the solid (LiBH4) and pass chlorine gass over it and lead the gas that will be produced over lithium with heating. Then dissolve it in water, the insoluble material should be boron. This is harder then the idea of passing the trimethyl borate over lithium but I'm not sure that idea will work. My references are from Wiki. Also thanks to Cheddite Cheese for help with the reaction.

[Edited on 23-10-2013 by bismuthate]

[Edited on 23-10-2013 by bismuthate]

WGTR - 23-10-2013 at 16:56

Don't worry, no offense taken. It's not the first time I've made bad assumptions, or outright hair-brained conclusions. I promise I'll even make some more.

After a bit more reading, since H3BO3 is barely disassociated in water anyway, it's likely that excess HCl won't affect its solubility very much. That was one bad assumption that I made.

The chart for the H3BO3 is actually in g/100mL, so everything has to be multiplied by 10 to obtain g/L.

Here's what was bugging me. For the equation:

Na2B4O7 + 2HCl + 10H2O <=> 4H3BO3 + 2NaCl + 5H2O

...I saw a possible issue. At 20C, borax solubility looks like 27 g/L (0.134 mol/L). For every mole of borax, 4 moles of boric acid reside on the other side of the equation. The solubility of boric acid at 20C is 47.2 g/L (0.763 mol/L).

(0.134 mol Na2B4O7/L) * (4 mol H3BO3/ 1 mol Na2B4O7) = (0.536 mol H3BO3/L)

Just going on solubility of the two salts alone, it looks like the equilibrium of the equation sits almost right in the middle.

Does a strong acid actually bring a fairly insoluble salt (with strong cation, weak uncommon anion) into solution, essentially pushing the reaction towards completion? Duh...I guess it does. Carbonates, hydroxides, sulfides, phosphates, etc. In the case of carbonates, the CO2 even leaves solution due to its limited solubility.

Ok, now I have a new project (like I don't have enough): Figure out how to calculate the equilibrium of a system like this. It bugs me that I can't do it off the top of my head.

So...nothing to see here...just somebody smoking their socks. Move along, move along now everyone.

violet sin - 23-10-2013 at 17:41

I enjoy reading and re-learning this stuff. Been about or greater than a decade since I was taking chem in college. The info is still there but deff not off the top of my head. And I certainly am not the type of person who derives joy from making others feel stupid, just FYI. I WAS good at equilibria calculations, but not any more... Sadly. Would make some of this miche easier. I plan on trying my suggestions as soon as I get back home. Still have a lineup of other projects waiting to happen too. But it's always nice to know if ones suggestions are as cut and dry as one assumes they are. Seems I have outlined a plan in checking the numbers for Upsilon, time for some action instead of words. I might even be able to pull this off here( 3.5 hrs from home/lab, working) we shall see soon enough. Thanks for being a sport about the criticism
Gotta love the good 'ole mason jar experiments. Check back in later when I know how it went.
So I found some borax and muriatic acid. Basically halved the 100g borax plan. 50g borax dissolved in ~120ml boiling water, added ~26ml HCl and now it's cooling outside. Hot water + pool acid = not fun in a small trailer. After a while ill transfer the jar to the freezer n see what comes of it. Not going to be able to weigh it till tomorrow but visual confirmation should be easy enough tonight. Prob try a flame test to see how brightly the sodium affects it for purity test, best I can think of right now(w/o additional chems)

[Edited on 24-10-2013 by violet sin]

Upsilon - 23-10-2013 at 20:06

I am still looking to extract boron from boric oxide without reduction with magnesium, but I need a suitable solvent to dissolve boric oxide without reacting with it (since adding water to boric oxide reverts it back to boric acid). Would acetone or isopropyl alcohol do the trick? I want to use aluminum foil to replace boron in single replacement, but I don't know if acetone/IPA will react with any boron formed.

violet sin - 23-10-2013 at 21:21

I started reading this-
http://library.iyte.edu.tr/tezler/master/kimya/T000628.pdf
On page 20, near the top it starts describing some methods for elemental boron. Said the most common was magnseiothermic reduction. ~92% pure and easy to dissolve the magnesium oxide away. Guess crystallized boron can withstand boiling HCl quite well. He(names Yelda,not sure boy or girl) then goes on to other methods. It's a thesis paper, and he intended to reduce it with aluminum also. Worth a look.

After sitting in the cold for a while I had a nice crop of crystals in the jar. Part were initiated on the liquid surface and fell down as a whole. So if you turn it just right the collective face reflects light from under the surface. Kinda cool looking. Would add a pic but no way to edit pics to proper size.

Stupid auto correct
[Edited on 24-10-2013 by violet sin]

[Edited on 24-10-2013 by violet sin]

violet sin - 25-10-2013 at 22:34

found some other papers, I know you didn't wanna do magnesium reduction, but some of these are kinda cool. each with their own difficulties. but one uses borax directly and the other works in a similar fashion but uses the oxide, making BC4 and BN respectively. not your intended product but I found little else on elemental boron, that wasn't covered in the link one post up. maybe fun, or useless to you

Synthesis of B4C from Na2B4O7+Mg+C by SHS Method
http://iopscience.iop.org/1757-899X/18/7/072007/pdf/1757-899...
they use high pressure argon to prevent magnesium evaporation. think it also helped efficiency. but you could probably add more magnesium.

Polycrystalline boron nitride (BN) powders were prepared by magnesiothermic reduction of boron oxide (B2O3) and ammonium chloride (NH4Cl) catalyzed with ferric oxide (Fe2O3) at temperatures ranging from 700 to 850 ℃ under ambient pressure.
http://www.cjcu.jlu.edu.cn/EN/abstract/abstract12540.shtml
I used google translate, also i tried to download the PDF from original page, but got something about ZnO instead of the BN paper?

Upsilon - 26-10-2013 at 20:29

I may just end up doing it with aluminum reduction. Boric oxide isn't too easy to produce in quantities large enough to blindly experiment in, and I just don't want to spend the effort creating the amounts needed to experiment with various solvents. Aluminum powder is dirt cheap and can be used in most reduction reactions, and it won't oxidize nearly as quickly as magnesium powder. The aluminum oxide formed can also be easily dissolved by some HCl.

Upsilon - 3-11-2013 at 18:34

From my first experiment, I ended up with a pitiful yield of boron trioxide, no more than a gram or so. I feel like the majority of this loss was in the reaction between the HCl and the borax. I started the experiment again with 200g of borax and I added excess HCl and stirred vigorously. I am trying to conserved as much material as possible, including that of which is dissolved in the HCl, so I am evaporating off the HCl. Just letting it sit outside was taking way too long, plus it rained when I wasn't home so I ended up with 2x the liquid I had to begin with. I constructed a small evaporating oven out of wood with an interior lined with aluminum foil. It uses two 60W incandescent bulbs as a heat source. It gets quite warm, to where I can see white HCl fumes being driven off, but the evaporation of the liquid is still taking quite some time but progress is at least noticeable.

Anyway, after it is mostly dry I am going to add HCl once again to ensure all borax has been reacted with, then I'll evaporate it off again. Afterward, I'll do some basic solubility distillation to hopefully get a substantial amount of boric acid.

eidolonicaurum - 3-1-2014 at 03:04

This works, I have done it myself.
Required: Borax, hydrochloric acid, aluminium or magnesium powder, mixture of potassium permanganate, aluminium powder and sulphur (3:2:1), crucible, source of heat
Dissolve the borax in the hydrochloric acid, and heat for about an hour depending on the quantity of reagent. Allow to cool, and then place in fridge or freezer to cool further. This precipitates out boric acid, which is much less soluble at lower temperatures. Filter this off. Dry, and heat to drive off water to form boric oxide. This will require a roaring Bunsen flame and a lot of time. Allow to cool, and grind up the resultant solid. Mix with the relevant quantity of aluminium powder, and ignite with the mixture. Dissolve the resultant solid in hydrochloric acid, and filter off the boron. This can be further purified and melted in an electric arc, if one can be generated.

Tdep - 3-1-2014 at 03:54

I actually happened to try a aluminium reduction of boron trioxide today, a ratio of 1.29:1 of oxide to aluminium powder. However, it failed to light, despite my constant attempts with Mg ribbon and chlorates (a method that lights most thermites i've found).

Anything I'm doing wrong or is it just being stubborn? It should ignite and the reaction should be sustaining enough to continue throughout the mass? (it was about 50g of reaction mixture)

Zyklon-A - 3-1-2014 at 12:02

You'll need a very fine powder, how small was the partical size?

Zyklon-A - 3-1-2014 at 12:09

Is it possible to isolate boron with aqueous reduction, (from Al or something), instead of a thermite? I have a lot of B(OH)3 and B2O3.
Edit: Sorry if this has been answered already, I haven't had much time to look for it...
Good yeilds are not high priority, I just want a gram or two for an element collection. I do have Al powder, but a a thermite reaction won't give high purity products right?

[Edited on 3-1-2014 by Zyklonb]

blogfast25 - 3-1-2014 at 13:45

Z:

There are no easy ways to produce boron. For an element collection, dissolve a few neodymium magnets (Nd2Fe14B) in strong hydrochloric acid. The small amount of insoluble black/brown residue is elemental boron.

Zyklon-A - 3-1-2014 at 14:18

How strong does the HCl have to be? I only have about 8-9 molar HCl.

Tdep - 3-1-2014 at 17:43

Quote: Originally posted by Zyklonb

You'll need a very fine powder, how small was the partical size?

400 Mesh aluminium powder, and the boron oxide was ground up pretty well, but I guess I can grind it more and have another go

Zyklon-A - 3-1-2014 at 21:26

No, that should easly be good enough, um, I have no idea why it didn't light, Mg ribbon has worked for all my thermates... I haven't tried this one though.

eidolonicaurum - 4-1-2014 at 01:14

You are right about the difficulty in lighting, I had the same problem. Try using magnesium powder, and if that doesnt light, burning the whole surface with a blow torch. Then, it does react, but not in a traditional thermite reaction. My pile ended up white all over.

Tdep - 4-1-2014 at 02:30

The ratio is 1.29:1? I'm not certain of my Stoichiometry here....

blogfast25 - 4-1-2014 at 07:10

Quote: Originally posted by Tdep

The ratio is 1.29:1? I'm not certain of my Stoichiometry here....

There's a procedure in Brauer's Preparative Inorganic Chemistry, see library.

http://library.sciencemadness.org/library/index.html

"According to Moissan, very impure amorphous boron, containing
about 80-90% B, is obtained by the reaction of B3O3 with magnesium.
According to Kroll the optimum yields are obtained as follows: A
fireclay crucible, approximately 20 cm. high and 16 cm. in diameter,
is painted with a paste of ignited MgO and sintered MgCl2 and
dried in a low-temperature oven. A mixture of 110 g. of B2O3,
115 g. of Mg shavings (the use of Mg powder frequently leads to
explosive reactions) and 94 g. of powdered S is placed in the crucible.
The reaction is started with an ignition pellet, after which it
proceeds vigorously. After the mixture has cooled, it is extracted
in water and then in dilute HCl for a week. The residue is treated
several times by heating with HF and HCl, washed with water and
dried in vacuum at 100°C. The yields are variable, with a maximum
of 46%."

Without sulphur that mixture is probably impossible to light and will not generate enough heat to self-sustain. Note that the author warns against explosive reactions. If you want to use Mg powder try on a small scale first.

For an ignition pill, I'd recommend a small amount of stoichiom. KClO3 + Al mixure, lit with a piece of Mg ribbon. Stand well back after lighting the Mg ribbon!

Washing the product will generate substantial amounts of H2S, so be prepared!

[Edited on 4-1-2014 by blogfast25]

Edit:

The formulation 100 g B2O3/115 g Mg/94 g S is strange though. It corresponds to 1.58 moles B2O3/4.73 moles Mg and 2.93 moles S. But acc. B2O3 + 3 Mg === > 2 B + 3 MgO, 1.58 moles of B2O3 require 4.74 moles Mg, leaving no Mg for the 2.93 moles of sulphur! It could explain why low yields of B are reported.

Also, acc. NIST values of Standard Enthalpy of Formation for B2O3 ( -1273.5 kJ/mole) and MgO ( - 601.6 kJ/mole), the reduction reaction B2O3 + 3 Mg === > 2 B + 3 MgO would have a Standard Enthalpy of Reaction of – 531.3 kJ/mole (of B2O3), which I would have thought to be more or less enough to make it self-sustaining, assuming you can get it to light.

If anything, I would adjust this formulation for a small scale test by reducing the amount of sulphur to 1 mole and increasing the level of Mg to 5.74 moles. The extra heat output from Mg + S == > MgS would make the total reaction heat about – 900 kJ/mole (of B2O3), similar to Classic Thermite.

[Edited on 4-1-2014 by blogfast25]

eidolonicaurum - 5-1-2014 at 11:22

Quote: Originally posted by Tdep

I actually happened to try a aluminium reduction of boron trioxide today, a ratio of 1.29:1 of oxide to aluminium powder. However, it failed to light, despite my constant attempts with Mg ribbon and chlorates (a method that lights most thermites i've found).

Anything I'm doing wrong or is it just being stubborn? It should ignite and the reaction should be sustaining enough to continue throughout the mass? (it was about 50g of reaction mixture)

Having tried to do this myself, it is almost impossible to light an aluminium/boric oxide mixture easily. I gave up on aluminium and used magnesium instead. It worked that time. I then chucked the products into dilute acid, which dissolves away everything except the boron. I got an amorphous brown powder.

eidolonicaurum - 5-1-2014 at 11:23

Quote: Originally posted by eidolonicaurum

You are right about the difficulty in lighting, I had the same problem. Try using magnesium powder, and if that doesnt light, burning the whole surface with a blow torch. Then, it does react, but not in a traditional thermite reaction. My pile ended up white all over.

Identical here. That, I think, is the way to do it.

MrHomeScientist - 7-1-2014 at 07:12

I'm almost certain I posted this earlier in the thread, but I don't want to go searching so here it is anyway.

I prepared elemental boron starting from boric acid roach killer and magnesium powder. I heated the boric acid on a clay dish until it melted and steam stopped coming off, to form boron trioxide. I ground this up (NOT an easy process, it's super hard) and mixed with a stoichiometric amount of magnesium powder. I ignited the thermite, and covered the reaction after it was finished to protect it from air while it cooled. I then broke up the 'cake' of products and digested in hydrochloric acid. Everything except the elemental boron will dissolve.

It's best to avoid aluminum in this reaction because it produces some aluminum boride, which is very inert and hard to separate from your wanted product. The magnesium analog (Mg₃B₂) reacts away to borane in the acid digestion step.

A link to my video on the process: http://www.youtube.com/watch?v=0QBCyOrjR2o

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