GammaFunction
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Low temperature Castner process for sodium?
I used the search engine, and I hope this isn't a repost.
Wikipedia says, without giving a reference, that low-melting sodium hydroxide monohydrate can be electrolyzed to give metallic sodium, at just
over 100C.
This inspired the idea of this simple apparatus to make sodium at just over the boiling point of water, by shielding the metallic sodium and hydroxide under mineral oil.
Fuming should be reduced or eliminated, and no hazardous pressure could build up.
Glass components wold be nice to observe the process, but I'm not sure how long it would withstand the NaOH even at this low temperature. Even some
form of plastic could conceivably work. The mineral oil itself will serve as a top window.
Any thoughts?
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elementcollector1
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"Hot electrochemical sodium" - stickied in Technochemistry, I suggest you read there!
Glass would withstand molten NaOH for about 5 seconds, and it would turn opaque in 1. No-go.
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GammaFunction
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Quote: Originally posted by elementcollector1  | "Hot electrochemical sodium" - stickied in Technochemistry, I suggest you read there!
Glass would withstand molten NaOH for about 5 seconds, and it would turn opaque in 1. No-go. |
Guess I didn't read this section. Oh well.
However, this thread seems to be devoted mainly to the high temperature Caster process, not this unusual lower temperature variant, which uses the
monohydrate to work just over the BP of water.
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BromicAcid
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Attached are pictures of a similar attempt from me. Starting with the most concentrated solution of sodium hydroxide I could make I began
electrolysis on a watch glass. The electrolysis creates plenty of heat from resistance of the solution which drives off water, and it also splits the
water to its elements driving it off as well. Eventually it gets to the hydrate. I do find it strange that sodium would form and be isolated from
the hydrate. Usually with the high temperature Castner cell the inital bit of electrolysis is to remove the water, no sodium is formed. Still,
stranger things have happened and I would chalk this up to overpotental. Nevertheless, without careful control of the temperature and current you
will overshoot this window and end up with anhydrous material.
In my attempt water continued to be split off as the mixture heated from the resistance until I was electrolyzing molten NaOH. You see several little
drops of sodium forming at the electrode. Afterwards there was a nice scorch mark on the picnic table due to the heat and the watch glass was nicely
etched. Would be a neat demonstration save the fine particulates of sodium hydroxide that end up floating around in the air.
Edit: Traced the Wiki source back to August 12, 2010 edit that has made it through revisions to now. I assume however you did the same as someone
asked the person who made the edit for a reference within the last few weeks. His response that "it is more of a proof of concept, rather than a
source of sodium" rings true considering the reaction simultaneously produces 4 mols of sodium with 6 mols of water and one mol of O2.
[Edited on 4/29/2013 by BromicAcid]
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GammaFunction
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| Quote: Originally posted by BromicAcid | | Attached are pictures of a similar attempt from me. Starting with the most concentrated solution of sodium hydroxide I could make I began
electrolysis on a watch glass. .... |
So it looks like you never made sodium from the monohydrate, but conditions were not carefully controlled.
The appeal of this process (if it works) is that the conditions are moderate, and an oil layer can be used to contain it (and any NaOH mist), and an
oil bath can be used to moderate the temperature.
Could NaOH hold its water of hydration so strongly that the monohydrate doesn't react with sodium (or reacts very slowly)?
The process seems to depend on the following two assumptions: 1) When Na is formed at the cathode, the liberated water of hydration doesn't react
with the sodium much before escaping as a gas; 2) the water in the remaining monohydrate doesn't react with the freshly created sodium.
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m1tanker78
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| Quote: Originally posted by GammaFunction |
The appeal of this process (if it works) is that the conditions are moderate, and an oil layer can be used to contain it (and any NaOH mist), and an
oil bath can be used to moderate the temperature.
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Hot NaOH will turn your oil to a milky white substance in no time. I believe the product is soap but I never got around to researching it.
Most likely, in Bromic's case at least, the large current density (small surface area of electrodes immersed) caused enough local heating to dehydrate
the lye around and between the electrodes. By the looks of the burnt table, it got quite hot..
Tank
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BromicAcid
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| Quote: Originally posted by m1tanker78 | | Most likely, in Bromic's case at least, the large current density (small surface area of electrodes immersed) caused enough local heating to dehydrate
the lye around and between the electrodes. By the looks of the burnt table, it got quite hot.. |
Indeed, my point was that one could start out with a hydrated form and then through the electrolysis and heat and resistance end up electrolyzing
molten hydroxide and not stop at the hydrate. To do what you are proposing will likely involve the ability to cool the setup appropriately once at
temperature to keep it in range or else it will quickly turn into a traditional Castner cell.
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GammaFunction
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| Quote: Originally posted by m1tanker78 |
Hot NaOH will turn your oil to a milky white substance in no time. I believe the product is soap but I never got around to researching it.
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Is this true even for a mineral oil, rather than an organic oil? Mineral oil should not saponify. I can't find any evidence that hydrocarbons react
with hydroxides.
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elementcollector1
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It's not true in my experience - KOH heated to 200+ C with mineral oil and kerosene had no reaction that I could discern (at least, nothing forming a
milky white substance).
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ScienceSquirrel
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If your mineral oil or kerosene is saturated hydrocarbons then it will not react with potassium hydroxide under any conditions that you are likely to
be able to achieve in a home lab.
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Fantasma4500
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IIRC 316 steel can withstand hydroxides
using it as cathode might work aswell as heating element / container
but i believe this is more fitting for technochemistry anyhow
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blogfast25
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Where precisely does it say that? If this is indeed possible, countless experimenters would have reported on it.
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GammaFunction
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It has been edited out.
It seems vaguely plausible only if the sodium were at least fractionally protected against the water of hydration. Maybe in the immediate
neighborhood of the cathode, the current density and heat are high enough to perform the anhydrous reaction.
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blogfast25
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Gamma:
Perhaps. It'll be like walking a tight rope, I think. This kind of thing could be OK for demo purposes (making micro quantities of Na) but I doubt if
it can be reasonably scaled up.
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kristofvagyok
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Just a thought. If you electrolyze NaOH.H2O to get sodium, it will immediately react with the
water from the monohydrate to form back the NaOH and producing some H2. It would be a really expensive and hard way to produce anhydrous NaOH 
P.S.: Anhydrous NaOH/KOH could be made by melting it in a silver pot till it boils (the water comes out). And be careful, because it eats everything.
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GammaFunction
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| Quote: Originally posted by kristofvagyok | Just a thought. If you electrolyze NaOH.H2O to get sodium, it will immediately react with the water from the monohydrate to form back the NaOH and
producing some H2. It would be a really expensive and hard way to produce anhydrous NaOH
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Yes. That's the most evident thing that would happen. It might work if the immediate zone around the cathode were hotter and water-depleted
(because the current density is highest there). One could imagine a partly efficient process in which some of the sodium reverted back to the
hydroxide, and some water was forced out of the system where it didn't react with the sodium.
So far, there was exactly one unsupported reference to this process (wikipedia), and it has been edited out in the past day, probably because this
thread drew attention from a justified skeptic.
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elementcollector1
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I have heard that if you take regular NaOH, stick in your electrodes (w/o melting yet), turn on the power, and add a drop of water or two, ohmic
heating will do the rest - no heat required!
I have not tested this yet, but it would be interesting to try someday.
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