UncleJoe1985
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getting sodium metal from aqueous electrolysis?
Although it seems like a dumb question, I'm wondering if the electrolysis of sodium salts in water can yield sodium metal hypothetically given the
right conditions?
Obviously, there are 2 factors working against it:
1. Reduction of H+ is favored over reduction of Na+
My understanding is H has a bigger ionization energy (1312 kJ/mol) than Na (496 kJ/mol) meaning H+ has more electrostatic energy. So by the principle
of natural affinity for the lowest potential energy state, the electron will go to hydrogen.
2. Immediate reduction of any produced Na by water
Contrast this with propylene carbonate, which doesn't steal electrons (aprotic)
Can both of these barriers be overcome by higher voltages? I know when there are multiple competing reactions like Cu+2 + 2e-
-> Cu and 2H+ + 2e- -> H2 at the cathode, a high current density will cause even the less spontaneous reaction
(H+ reduction) to start happening.
Likewise, does cathodic protection prevent the sodium from being reduced by water?
I don't want to try this. I'm just asking theoretically can it be done?
[Edited on 3-6-2015 by UncleJoe1985]
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j_sum1
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Short answer.
No.
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j_sum1
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Longer answer.
You actually appear to understand the reasons pretty well.
As you say, H2 production is favoured over Na and so that is what you get.
Even if you did happen to produce any Na, it would immediately react with the water to produce the Na+ you started with as well as H2.
So, no. It can't be done. Believe me: if it was possible, then we would all be doing it. (Actually, if it was possible then sodium wouldn't be
nearly as exciting or useful a metal.)
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WGTR
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It partially depends on how much water you need for something to be called "aqueous".
Molten sodium hydroxide can contain some moisture, yet it will still yield sodium metal at high enough current densities. At low current density the
electrolysis of water predominates until it dries out. Even then, water is produced when this salt is electrolyzed, normally limiting efficiency to
< 50%.
At room temperature, no, unless you include electrolysis at a mercury cathode.
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Corrosive Joeseph
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http://www.sciencemadness.org/talk/viewthread.php?tid=9797#p...
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nezza
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It can be done via a roundabout route. Because of the high overpotential at a mercury cathode electrolysing aqueous sodium salts with a mercury
cathode will produce a sodium amalgam. The amalgam is then run into pure water and reacts giving sodium hydroxide. This method has been used
commercially. By extension I presume the mercury could be distilled off to leave sodium. Distilling mercury would obviously be extremely hazardous and
require precautions.
If you're not part of the solution, you're part of the precipitate.
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morganbw
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Quote: Originally posted by nezza  | | It can be done via a roundabout route. Because of the high overpotential at a mercury cathode electrolysing aqueous sodium salts with a mercury
cathode will produce a sodium amalgam. The amalgam is then run into pure water and reacts giving sodium hydroxide. This method has been used
commercially. By extension I presume the mercury could be distilled off to leave sodium. Distilling mercury would obviously be extremely hazardous and
require precautions. |
I should search and reference this but I do not have the time this morning.
Mercury/sodium amalgam goes solid with a very small % of sodium in the mercury. I am sure it is too small an amount of sodium to make this feasible
for an amateur.
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morganbw
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From WIKI
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Structure and compositions[edit]
No particular formula is assigned to "sodium amalgam." Na5Hg8 and Na3Hg are well defined compounds. In sodium amalgams, the Hg-Hg distances are
expanded to around 5 Å vs. about 3 Å for mercury itself.[1] Usually amalgams are classified on the weight percent of sodium. Amalgams with 2% Na are
solids at room temperature, whereas some more dilute amalgams remain liquid.[2]
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Just saying, much mercury/little sodium.
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UncleJoe1985
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| Quote: | | electrolysis at a mercury cathode... The amalgam is then run into pure water and reacts giving sodium hydroxide |
How's that an advantage to just using a diaphragm cell for making NaOH?
That's quite a curiosity, but the crux of my question is does the excess supply of electrons on the cathode (cathodic protection) prevent any sodium
formed from being reoxidized?
Actually, my question does have a more practical purpose or else I wouldn't be on this site. I saw some videos on making sulfuric acid from a solution
of copper (II) sulfate. The copper gets plated onto the cathode, leaving sulfuric acid in solution. It seems cathodic protection prevents the copper
from being dissolved by the acid.
I immediately thought, how wonderful it would be if that same method could be used to make nitric acid from a nitrate solution!
But I'm doubting it will work because nitric acid brutally attacks copper unlike sulfuric acid. Ideally, I'd also want to use easier to obtain
nitrates like KNO3 instead of Cu(NO3)2 which will plate out even more reactive metals.
Maybe I should've asked that question directly, but I thought electrolysis of aqueous salts would be more well known.
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WGTR
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| Quote: Originally posted by UncleJoe1985 | | Quote: | | electrolysis at a mercury cathode... The amalgam is then run into pure water and reacts giving sodium hydroxide |
How's that an advantage to just using a diaphragm cell for making NaOH?
That's quite a curiosity, but the crux of my question is does the excess supply of electrons on the cathode (cathodic protection) prevent any sodium
formed from being reoxidized?
Actually, my question does have a more practical purpose or else I wouldn't be on this site. I saw some videos on making sulfuric acid from a solution
of copper (II) sulfate. The copper gets plated onto the cathode, leaving sulfuric acid in solution. It seems cathodic protection prevents the copper
from being dissolved by the acid.
I immediately thought, how wonderful it would be if that same method could be used to make nitric acid from a nitrate solution!
But I'm doubting it will work because nitric acid brutally attacks copper unlike sulfuric acid. Ideally, I'd also want to use easier to obtain
nitrates like KNO3 instead of Cu(NO3)2 which will plate out even more reactive metals.
Maybe I should've asked that question directly, but I thought electrolysis of aqueous salts would be more well known. |
I'll answer some of this below:
There once was a good video (Youtube, I think?) that showed a very slick mercury cell demonstration. I can't find it for whatever reason. My
understanding is that the sodium amalgam is cathodically protected up to a certain small percentage, above which hydrogen production becomes the
dominant reaction.
Your particular example with nitric acid wouldn't work, because a copper cathode happens to be a very efficient catalyst for nitrate ion reduction.
Using a platinum anode and a copper cathode, I've reduced dilute nitric acid all the way to ammonia. You can actually watch the pH rise as time goes
on.
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Marvin
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Cathodic protection works well for metals that can be electroplated.
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