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Quantum
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patu seems to have a good small scale idea for sodium. However then Tacho got hot NaOH on his person; thay does not sound very fun. I just want a few
beads of Na nothing fancy like you guys at least not yet.
Has anyone tried patu's method besides Tacho and can they provide more detail?
What if, what is isn\'t true?
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Tacho
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Bromic acid's method seems the easiest so far. If an iron (steel) wire can be used as electrode it would be even more simple.
About Patu's way, I must say that my lab techniques are sloppy and hurried, and I used Cu wire as the loop electrode.
For what I have seen so far, though, electrolisys of molten NaOH is an intrinsically messy procedure. It pops and fizzes a lot.
Edit: A try of Bromic's method using a small iron loop as electrode would be a nice experiment.
[Edited on 30-4-2004 by Tacho]
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patu
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I must admit that its does pop the newly made sodium out of the crucible. the cool thing is, is that it calms down greatly after about a minute and a
half. After that the molten sodium quietly forms waiting to be dunked in the mineral oil. For the more reactive part of it i'm suited up in a
heavy coat and a welder's mask and gloves.
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Quantum
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Patu! You realy should post more; you seem to have lots of interesting ideas. Perhaps you just read 99% of the time.
Thanks for the information. I think I would just stay far away for the first few mins!
Edit: Are you running it on a hot plate or a flame? Do you hold it just above the mp of NaOH or much hotter?
[Edited on 1-5-2004 by Quantum]
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The_Davster
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When I made my bit of sodium I found that it was best to melt the sodium hydroxide over a burner, then once molten remove the crucible from over the
burner and quickly place the electrodes in and turn on the power, the current kept the area between the electrodes molten. When I was done I decided
to see what would happen if I heated it again. In under a minit the sodium blob had dissolved into the sodium hydroxide.
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patu
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I use a hot plate when I make sodium. I dont even own a bunsen burner. the methods I use are very crude yet very effective. I have little bottles full
of sodium and sodium/potassium alloys thet I have made.
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Strepta
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BromicAcid's use of resistance heating
I've been following this topic for awhile and have tried several experiments. Just this evening I tried Bromic's experiment using a watch
glass, OTC Red Devil Lye and two electrodes made of Ni 200 sheet metal. I started the reaction using a propane torch then turned on the power which
was an 0-14v, 0-5 amp bench power supply. Spacing between electrodes was approx. 1cm, depth a few mm. the initial curent was 1-2 amps (fluctuated a
lot) and stedily increased to about 4 amps as the H2O was electrolyzed out. In a few minutes the sparks began to appear, then the first shiny
globules of Na. My experience was almost identical to that described by Bromic. As to chemoleos questions, I suggest that the phase change
(solid-liquid) interface which is enabled by resistance heating keeps the temperature right at the melting point, not allowing it to climb to the
point where the Na becomes soluble forming the metalloid. Also, all molten material is in the vicinity of the electrodes--becoming solid 1-2 cm
away--this protects the glass from attack by the NaOH. All in all a very impressive little demo made more so by the minimum of fixturing required.
Haven't yet found a very effective way to remove the Na; may try the wire loop suggested by PATU. More later...
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Quantum
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Look! For I have found good infomation
SAS Chemistry Guide to making small amounts of Sodium
They decribe in good detail the proccess that patu uses . They use a burner instead of a plate but they even use a 'tin' can. Worth reading
to all interested in making small amounts of sodium.
Im going to try it tommorow!
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Strepta
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Bromic's technique with patu's wire loop
I replaced the nickel sheet metal cathode with a loop of iron wire (.35 cm dia.). This allowed me to capture almost all visible Na.  I hold the loop with a gloved hand--yes, molten NaOH is nasty-- and work it around the
melt to get the Na inside. I then lift it out and slap it against the test tube containing motor oil (just because it was nearby). I ran the setup
this way for an hour, removing the little balls of Na and quickly returning the loop to the melt. I was usually fast enough to get it back before the
NaOH solidified. I kept the propane torch on low flame nearby so if I was too slow, I could re-melt and be making Na within a few seconds. this
technique seems to work acceptably well (at least for me) to make small amounts of Na. I'm still amazed that a mere 35 watts in ( I'm
running this setup at 10V and 3.5 A) is enough to keep the melt at 315C without any attempt to insulate from ambient. I have not yet perfected a way to clean up the Na that is produced. I've
melted it under paraffin, xylene and ordinary paint thinner, all three of which result in shiny balls of Na but there is still come contamination
(NaOH) visible each time. Ideas welcome.
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BromicAcid
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| Quote: |
I've melted it under paraffin, xylene and ordinary paint thinner, all three of which result in shiny balls of Na but there is still come
contamination (NaOH) visible each time. Ideas welcome.
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You need an inert liquid with a density of 1 g/cm3 or above with a boiling point of greater then 100C. This is the technique used to purify lithium
that I mentioned before, you melt the impure NaOH/Na mixture at the bottom and as it melts the pure-er Na will float to the top leaving the NaOH
behind. You could use something like nitrobenzene possibly. I think most chlorinated hydrocarbons will react with Na at these temps and that is the
massive drawback because most organic liquids that have a density of greater then 1 are halogenated hydrocarbons.
[Edited on 5/9/2004 by BromicAcid]
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Quantum
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Chemical burns!
I tried the sodium making using a spoon on my hotplate. After it melted I stuck a 9v transformer wires into it but nothing happened. I was convenced
that it was dead so I put the 2 wires on my tough to see if it would 'tingle'. Sadly the wires had a little bit of NaOH on them! I had to wash my mouth out and now I have a little burn. I used a diffrent PSU(5v
20a) and I saw a little sodium but I could not collect it. I think its redisolving in the NaOH even though I turned off the plate and lifted the spoon
off it when I applied the current.
In a way I failed but I did see a speck of shiny Na!
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Saerynide
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You freak!!! Dont lick live wires, no matter what the voltage is!!
To test if a PSU is working just shove the wires into tap water and see if it bubbles 
[Edit:] Im sure you know this already, but you should try to *show* it
[Edited on 9-5-2004 by Saerynide]
"Microsoft reserves the right at all times to monitor communications on the Service and disclose any information Microsoft deems necessary to...
satisfy any applicable law, regulation or legal process"
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Quantum
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I don't know what I was thinking! The sad thing was that I had a multimeter
sitting on a bench 5 feet from me the whole time. I didn't get much sleep last night.
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Tacho
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| Quote: | Originally posted by Cyrus
I have gotten NaOH on my tongue too,
(snip) |
4!
I have used thin ss pots to melt NaOH, and used the pots as anodes (not where Na forms). That's about 10 to 20 minutes of exposure. They
withstand it, but were dark grey/black where exposed. I must say I could never make sodium by electrolysis.
I read somewere that type 304 ss will dissolve in molten NaOH. The information itself is probably useless, but the key idea is that there are
different ss, with different properties.
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hodges
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My Results
I recently bought a power supply that is capable of putting out either constant voltage or constant current. It will put out any voltage from 0-30V
and any current from 0-5A.
So I decided to try some NaOH electrolysis. I placed some NaOH in an empty metal food can. I used the can as the anode and an ordinary copper wire
as the anode. I slightly wet the NaOH and then turned on the power supply. I found that a current of about 1A, which was about 10V, was required to
keep the NaOH melted enough to conduct. So I set a current of 1.5A, which was about 14V. There was, of course, a lot of bubbling from the cathode,
and occasionally an orange sparkle would rise from the cathode into the air. Kind of like when you light a blowtorch and some of the gas in the air
from before you lit it burns above it. There were also occasional pops of exploding hydrogen, accompanied by a yellow flame. After a few minutes, I
could see the shiny sodium floating near the surface of the melted NaOH at the cathode. After another minute or so, there was a loud crack and two
orange glowing pieces of molten sodium shot into the air. After that, there were smaller cracks regularly at about 10 to 15 second intervals.
Apparently the sodium was shorting between the electrode and the can, resulting in the regular pops. At this point I turned off the power supply.
I waited for the NaOH to cool. But when it did, I was surprised that I didn't see nearly as much sodium as I expected. There were a few small
pieces inside the NaOH, which I chipped away, but not much. I added water to the can and got quite a bit of fizzing, though, along with large orange
sparks. Also there was a small amount of sodium left on the cathode wire, and when I put a drop of water on it there was quite a bit of fizzing.
At that point, I cleaned up the mess, and put away the power supply. But then after everything was put away, I noticed what appeared to be two drops
of foam on the side of the metal sink. These drops were slowly bubbling. I picked one up on the end of a popsicle stick, and sure enough it looked
like metal surrounded by the foam. I dropped the small one, about the size of a pinhead, into water and it fizzed. I then dropped the larger one,
about the size of a small BB, into the water. It skated back and forth on the surface of the water, gave off large orange sparks, and finally
"popped" away from the surface of the water when it was all nearly consumed. The biting smell of NaOH vapor in the air was noticable
nearby.
Later I tried a similar experiment using a glass container and steel wires for the electrodes. I found it much harder to maintain the current flow,
and it took over 20 volts just to get an amp of current. There were no explosions, but every few minutes the sodium would burn at the cathode, for
several seconds accompanied by a yellow light and a decrease in conductivity. Upon cooling, I had no sodium to speak of. The solid mass did not
react with water, and there was just a small amount of fizzing when putting the cathode wire into water.
However, I did notice something strange in the second experiment. The steel anode wire was eaten away considerably (about 1/4 inch of it was gone).
Near the anode, the NaOH had a dark purple color. I'm wondering if the steel has manganese in it, and KMnO4 was formed. I know manganates can
be formed from oxides of manganese in melted alkalis, but I thought these were usually green until acidified.
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BromicAcid
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| Quote: |
The steel anode wire was eaten away considerably (about 1/4 inch of it was gone). Near the anode, the NaOH had a dark purple color.
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Anodic oxidation of iron to ferrate, I've been experimenting with this lately, seems to work reasonably well, color changes are significant and
usually somewhat quick. Acidification will result in almost instataneous gas evolution from decomposition of ferrate. That is why I choose nickel as
my number one electrode material in hydroxides.
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hodges
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Would iron ferrate be purple though?
------------------------------------------------
YES. Look for the ferrate thread and keep this on topic please.
[Edited on 6-6-2004 by vulture]
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The_Davster
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Hodges; I have found it best to remove the sodium from the sodium hydroxide while it is still molten. Upon solidification, some of the sodium seems
to be absorbed into the sodium hydroxide.
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BromicAcid
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Sodium is much more soluble in cooler hydroxide then hot. For example, 25.3 g of Na will dissolve in 100 g of NaOH (l) @ 480C but at 800C only 6.9 g
will dissolve. The papers I have say the solubility increases consistently to the solidification point of the hydroxide melt and therefore it is more
economical to remove the Na before cooling the melt. (However if you run your melt too high you increase cell corrosion and decrease yeild by having
the sodium react with it's hydroxide so there is some give an take, however I believe it is the general consensus to run the cell at as low a
temperature possible.)
To remove your sodium you could try a chilled iron wire, upon touching the surface the sodium should freeze to it and be able to be scraped off, sorry
if I've mentioned this method before.
Some information relating to the actual electrochemical process and reason for low yeilds:
"The hydroxide is electrolyzed
NaOH ---> Na + OH-
the sodium appears at the cathode and, at the anode, the hydroxyl is resolved into water and oxygen. The hydrogen that appears at the cathode is the
product of a secondary chemical reaction, between the sodium at the cathode and water formed at the anode. It is therefore possible to have both
hydrogen and oxygen liberated at the anode, and explosions may result..... The sodium which diffuses to the anode may also react with the oxygen there
being evolved, forming sodium peroxide, and the later may react with more sodium forming the monoxide."
So wouldn't it be possible to add more water to the mix to react away the Na2O2 and such thereby increasing overall yields? Although adding H2O
to molten hydroxide would be a bad thing, commercial hydroxide has appreciable water in it so just adding some of that could add the necessary water
to hydrolyze the oxides of Na to the hydroxides and keep the yield high correct? However this would only be a matter of concern in large scale
projects.
(500th post, I am the first International Hazard!)
[Edited on 6/7/2004 by BromicAcid]
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Marvin
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Where does the sodium solubility data come from? Its not what I'd have expected. Having said that, it shouldnt really be relevent as to succeed
you have to be operating the cell only just above the melting point of the lye.
Either 7 or 12% of carbonate improves the yeild (I forget which and nolonger have access to the book).
The idea about adding water to react with the sodium peroxide forming at the anode is..... questionable. You said yourself water is forming at this
electrode anyway, and its a bad thing because it depletes the sodium produced by diffusion.
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BromicAcid
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The data came from the "Complete Treatise on Inorganic Chemistry" in the section about the isolation of the alkali metals. Additionally 1 -
3% NaCl is supposed to improve yields in a castner cell but cell corrosion increases dramatically every percent so there is a massive tradeoff.
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Esplosivo
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Sodium borate, commonly found as borax, seems to be as good as sodium ethanoate for molten electrolysis. It has a melting point of 75 deg Celcius, and
is said to be easily available, though I cannot find it still. The sodium formed would be in the solid state. Would this work out? Thanks
The MSDS of sodium borate
Edit: What would form at the anode of such a cell?
[Edited on 9-6-2004 by Esplosivo]
Theory guides, experiment decides.
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chemoleo
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Lol, that's a mistake that has been done before (see sodium acetate etc)
Your sodium borate contains crystal water, Na2B4O7 . 10H2O - and you can imagine what happens to the nascent sodium once it contacts water...
Never Stop to Begin, and Never Begin to Stop...
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Esplosivo
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Yes right, my bad. If may I ask, can't the sodium borate/acetate by dried, like for example heating in a microwave like other normal ionic salts.
I know excessive heating will result in decomposition, but will removing the water require heating to such high temp for the salt to decompose?
Edit: I've found out that sodium acetate loses the three water molecules from the trihydrate form at a temp of approx. 123 deg Celcius and
decomposes at a temp above 300 deg celcius. Two different MSDS sheets give different melting points of the anhydrous acetate. What is it really -
58deg Celcius or 324 deg celcius?
Edit 2: Sorry for the stupid question. Just found out that at such a temperature both the acetate and the water will be in the gaseous state and
therefore cannot be seperated.
[Edited on 9-6-2004 by Esplosivo]
[Edited on 9-6-2004 by Esplosivo]
Theory guides, experiment decides.
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hodges
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Usually when I see a low "melting" point for a hydrated salt I find that the "melting" is actually the water being released, which
the less-hydrated salt then dissolves in.
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