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AndersHoveland
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I have been thinking about the possibility of reducing sodium hydroxide with aluminum foil. So a calculation of the expected enthalpy of formation of
such a reaction may be helpful, to get some idea as to whether such a reaction would be expected to be favorable.
The enthalpy of formation for Al2O3 is -1669.8 kJ/mol, while the value for Na2O is -414.2 kJ/mol.
As Al2O3 contains 3 times as many oxygen atoms per mol, 3 times 414.2 equals 1242.6, which is still less than 1669.8, so aluminum has more affinity
for oxygen than sodium. And indeed an exothermic thermite reaction between sodium hydroxide and aluminum powder can produce sodium.
http://www.youtube.com/watch?v=908rjHQ5mmc
The enthalpy of formation for AlCl3 is -705.63 kJ/mol, while the value for NaCl is -411.12 kJ/mol.
As 3 times 411.12 equals 1233.36, sodium has more affinity for chlorine than aluminum. And indeed, the reduction of aluminum chloride by elemental
sodium was first done by H. Sainte-Claire Deville, although H. C. Ørsted had previously used potassium instead.
But of course the interaction with the alcohol would affect the enthalpy of formation, increasing the affinity of sodium for oxygen. A quick
estimation of this effect can be made by comparing the enthalpy of formation for sodium hydroxide, which is no doubt even more favorable than sodium
alkoxides (sodium alkoxides vigorously hydrolyse with water).
NaOH -425.93 kJ/mol
H2O -285.83 kJ/mol
Na2O is -414.2 kJ/mol
So the hydration of sodium oxide to anhydrous sodium hydroxide should release 151.83 kJ for each mole of Na2O reacted.
Na2O + H2O --> 2 NaOH
So it can be inferred that the presence of tert-butanol would not significantly affect the affinity of sodium for oxygen, meaning that the reduction
of a sodium alkoxide by aluminum should still be energetically favorable.
The competing affinities between sodium and aluminum for fluorine apparently is more complicated:
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It will be noted that when aluminum fluoride is in excess to that contained in cryolite (NaF)6Al2F6, aluminum does not reduce sodium fluoride, and on
the other hand, when sodium fluoride is in excess, aluminum does reduce sodium fluoride.
Metallurgical and Chemical engineering, Volume 11, p178 (1913)
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condennnsa
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Well aluminum powder does reduce NaOH. But only to NaH I think. This was discussed in the unconventional sodium thread I think. I have tried it
myself. ground up equal amounts of NaOH and Al powder , and started heating it with a torch. First it starts bubbling a lot of hydrogen gas, but this
is not the reaction, if you heat the pile to red-orange heat, at some point a thermite-like reaction occurs, and the whole thing gets white hot. IRC
some people successfully isolated potassium with this approach, but not sodium. After my pile cooled down i broke it up and could not see any shiny
metal, when i put it water it reacted releasing a lot of hydrogen , but not like sodium, not violent. So I concluded that aluminum is oxidised to
Al2O3, and what's left behind is mostly sodium hydride which reacted with water.
[Edited on 8-2-2012 by condennnsa]
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AndersHoveland
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That could be possible. KH is apparently a more reactive reducing agent than NaH. NaH is also more stable, decomposing to its elements at a much
higher temperature than KH. NaH begins to decompose at 330°, KH at 200°. ( O. Ruff ) Wikipedia gives much higher values at around 800° and
400°, respectively.
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Pyrrolidone is completely metallated in 2-3 hr at room temperature with KH, while no reaction is observed with either a potassium dispersion or NaH.
At elevated temperatures (above 75 C), both K and NaH react, slowly liberating hydrogen, however, at 75 C the amide evidently undergoes secondary
decompositions as fast as formed for the solutions never develop sufficient base strength to deprotonate triphenylmethane indicator detectably.
The solvents which appear most suitable for reactions with KH are ethers. Aromatic hydrocarbons are inert to KH.
"Potassium hydride, highly active new hydride reagent"
Charles Allan Brown
J. Org. Chem., 1974, 39 (26), pp 3913–3918
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However, an alternate explanation for why potassium was isolatable from the thermite mixture is that, at such high temperatures, it is significantly
more volatile (lower boiling point).
I do not know whether potential formation of NaH is possible in a liquid tert-butanol reaction using NaOH, since NaH reacts with alcohols to form
sodium alkoxides and H2.
Also, I found something which might possibly explain why only sterically hindered alcohols, such as tert-butanol, seem to work in the reaction.
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NaH has been successful only in reactions involving relatively acidic compounds (e.g. ethyl acetoacetate, unhindered alcohols )
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[Edited on 8-2-2012 by AndersHoveland]
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blogfast25
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Quote: Originally posted by AndersHoveland  | I have been thinking about the possibility of reducing sodium hydroxide with aluminum foil. So a calculation of the expected enthalpy of formation of
such a reaction may be helpful, to get some idea as to whether such a reaction would be expected to be favorable.
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It’s already been done, on this forum. It creates a mess, from which the sodium is practically irrecoverable.
To apply it as on this thread would require Al butoxide to form and be sufficiently stable.
Look at NaF + Al:
NaF: STP HoF = - 577 kJ/mol (Wolfram Alpha)
AlF3: STP HoF = - 1423 kJ/mol (NIST Webbook)
NaF + 1/3 Al == > Na + 1/3 AlF3: STP HoR = 577 + 1/3 x (- 1423) = + 102 kJ/mol
But: AlF3 MP = 1291 C, Na BP = 883 C. By heating a mixture of NaF and Al to about 1000 C in vacuum it may be possible to distil off sodium vapour,
'pulling'. the equilibrium to the right... I, for one, will not be trying this. 
I will 'shortly' (it's bitterly cold here) try dihydro myrcenol (a tertiary alcohol with longer chain length) for K synth.
[Edited on 8-2-2012 by blogfast25]
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TheChemINC
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I didnt understant the german version of this but i know that you can heat magnesium and potassium hydroxide in mineral oil while it is refluxing.
Then if you add a tertairy alcohol, like T-butanol for instance.... it will catalyse the reaction and help coalesce the potassium metal that is
formed..... check out a video by a guy named NurdRage on youtube.... most of you guys have probably heard of him, but he gives an excellent
explanation of the process. also, not to change the subject, but does anyone know where i could get or how to make t-butanol?
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bfesser
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If you had bothered to read the thread, you'd find that everything you've just stated had been discussed already. NurdRage even participated in the
discussion, if I remember correctly. As for t-butanol, refer to your cross-post--and use the search function next time. You'll find that most
members don't take too kindly to strangers who don't know how to use Google or the forum search function.
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TheChemINC
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oh, sorry. didnt mean to piss andyone off..... :/
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White Yeti
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Don't sweat it, we're just pointing you in the right direction. Many long time members are rough on beginners.
Welcome to the forum!
"Ja, Kalzium, das ist alles!" -Otto Loewi
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TheChemINC
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thanks...... i was just trying to figure this out 
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Waffles SS
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I didnt read all of this 37 pages( i just read 10 or less pages) but i am interest to know : does this method is usable for making lithium metal?
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White Yeti
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| Quote: Originally posted by Waffles SS | | I didnt read all of this 37 pages( i just read 10 or less pages) but i am interest to know : does this method is usable for making lithium metal?
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It should be, but considering it doesn't yet work with sodium, it would be quite a stretch. Besides, you can get lithium from batteries. This is not
the case with sodium or potassium metal.
"Ja, Kalzium, das ist alles!" -Otto Loewi
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TheChemINC
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you can use lithium as a reducing agent to make potassium or sodium though.....
i took lithium foil from a battery and poured potassium chloride over it. after i heated it up with a bunsen burner, i place a watchglass over the top
of the test tube i used. then i broke apart the test tube and threw the material that formed in the bottom into a beaker of water and purple flames
appeared instantly. it was pretty cool.... im not sure that it is efficient and the product is very crude and mixed with tons of other crap.....
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blogfast25
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| Quote: Originally posted by Waffles SS | | I didnt read all of this 37 pages( i just read 10 or less pages) but i am interest to know : does this method is usable for making lithium metal?
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The patent specifies Na (but to do harder than K), K and Cs. Definitely not Li.
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blogfast25
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| Quote: Originally posted by TheChemINC | you can use lithium as a reducing agent to make potassium or sodium though.....
i took lithium foil from a battery and poured potassium chloride over it. after i heated it up with a bunsen burner, i place a watchglass over the top
of the test tube i used. then i broke apart the test tube and threw the material that formed in the bottom into a beaker of water and purple flames
appeared instantly. it was pretty cool.... im not sure that it is efficient and the product is very crude and mixed with tons of other crap.....
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It's not efficient at all that way. Really for it to work you need to heat your KCl/Li mixture in vacuum and above the BP of potassium (it's lower in
partial vacuum than at atmospheric pressure, of course), thereby distilling off the potassium.
Someone here did this with Li/CsCl (search for it...)
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TheChemINC
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thats interesting..... i didnt know that it would work better in a vaccum. but i guess it would work better if it was void of O2 so it wouldnt
oxidise. but when i did it, it wasnt to collect any samples of potassium metal. im going to use the method that this thread explains. it was more just
for the fun of it, and to show my chem teachers....
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White Yeti
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| Quote: Originally posted by TheChemINC | | im going to use the method that this thread explains. it was more just for the fun of it, and to show my chem teachers.... |
Love the enthusiasm, but if you look at this thread, you will see that some people took MONTHS to get this to work. Albeit, the Raging Nurd made a
video on the subject, but it undermines all the time that went into getting this method to work.
In short, don't get your hopes up. If you do get it to work, be sure to post some pictures
"Ja, Kalzium, das ist alles!" -Otto Loewi
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Waffles SS
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Thanks everyone,
I found out this is impossible to get lithium metal by this method
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blogfast25
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| Quote: Originally posted by White Yeti | Love the enthusiasm, but if you look at this thread, you will see that some people took MONTHS to get this to work. Albeit, the Raging Nurd made a
video on the subject, but it undermines all the time that went into getting this method to work.
In short, don't get your hopes up. If you do get it to work, be sure to post some pictures |
Nonsense, Yeti: the 'trailblazers' took some time to get it right at first but for anyone capable of following simple instructions and in possession
of the required chemicals this is a walk in the park.
Bear in mind a 'K synthesis' takes 3 - 4 hours of simmering the reagent mix at 200 C to get those balls of K, that's all...
This reaction was considered for a long time to be mysterious, sensitive, 'Goldielocks like' etc etc but those days are long gone, trust me.
[Edited on 25-2-2012 by blogfast25]
[Edited on 25-2-2012 by blogfast25]
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MyNameIsUnnecessarilyLong
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Anyone here have experience with doing a large scale of this reaction? Something at least 10:1? My aim is really to do about 40:1
Is there an issue with contact area or temp regulation, and overall reaction time?
I recall people saying it works better to not stir the rxn much, but would something like an overhead stirrer be needed for larger scales?
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watson.fawkes
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If you've never scaled up a reaction
to pilot scale, I would heartily recommend this one not be your first one.
If you insist, however, start with the ΔG of the reaction, an estimated reaction rate, and the heat capacities of the materials.
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blogfast25
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Seconded totally: never scale up if you haven't bench tested it yet. Are you equiped to reflux 2 L of kerosene (or equivalent) for 2 to 4 hours at 200
C? Do you have fire fighting procedures in place should things go south? A large enough mantle heater and reflux cooler? 
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Pok
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I did it several times in 10:1 scale. It works. By the way, the patent also says "1 liter".
When I want more, I make two 1-liter-batches simultaneously. Yields about 70 grams in total.
Recently, it didn't work anymore. Some problems which I can't explain up to now. But if you were successful in making the small scale reaction,
1-liter-batches will also work. Don't use too fine Mg powder! Otherwise the hydrogen will evolve too fast and it can become extremely dangerous
(inflaming hydrogen/shellsol aerosol mixture), especially in larger scale! Also, only do a large scale reaction in a suitable laboratory with fire
safety equipment (against burning paraffins and burning alkali metals!).
Temperature control, reaction time and stirring method were the same as in small scale.
[Edited on 2-3-2012 by Pok]
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Neil
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| Quote: Originally posted by watson.fawkes | If you've never scaled up a reaction
to pilot scale, I would heartily recommend this one not be your first one.
If you insist, however, start with the ΔG of the reaction, an estimated reaction rate, and the heat capacities of the materials.
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Pffft what's the worse that could happen
<iframe sandbox width="420" height="315" src="http://www.youtube.com/embed/C561PCq5E1g" frameborder="0" allowfullscreen></iframe>
Oh... well other then that I mean...
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MyNameIsUnnecessarilyLong
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| Quote: Originally posted by blogfast25 |
Seconded totally: never scale up if you haven't bench tested it yet. Are you equiped to reflux 2 L of kerosene (or equivalent) for 2 to 4 hours at 200
C? Do you have fire fighting procedures in place should things go south? A large enough mantle heater and reflux cooler? |
Yes to all of those. But I am planning to reflux it in an expendable thick-bottomed erlenmeyer on hotplate instead of rbf in mantle. Last time I used
a rbf and the KOH severely etched the glass, in some places the glass was left less than a 1/16" thick.
I will have the set up in a heavy gauge steel cabinet. Any fire should be contained. The box also has large gaping holes with steel shutters on them
so any fuel explosions shouldn't blow it apart. If anything escapes, there is nothing within 40 feet around or above it that isn't metal or concrete.
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MyNameIsUnnecessarilyLong
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| Quote: Originally posted by Pok | I did it several times in 10:1 scale. It works. By the way, the patent also says "1 liter".
When I want more, I make two 1-liter-batches simultaneously. Yields about 70 grams in total.
Recently, it didn't work anymore. Some problems which I can't explain up to now. But if you were successful in making the small scale reaction,
1-liter-batches will also work. Don't use too fine Mg powder! Otherwise the hydrogen will evolve too fast and it can become extremely dangerous
(inflaming hydrogen/shellsol aerosol mixture), especially in larger scale! Also, only do a large scale reaction in a suitable laboratory with fire
safety equipment (against burning paraffins and burning alkali metals!).
Temperature control, reaction time and stirring method were the same as in small scale.
[Edited on 2-3-2012 by Pok] |
Thanks Pok. I will do a run at 1L first
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