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kristofvagyok
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Quote: Originally posted by APO  | | I would rather use carbonic acid to produce carbonates. I didn't look at the sticky thread so much because I had trouble finding dense solvents, can I
get an example of a solvent at least twice as dense as water? |
Just a suggestion: learn some chemistry, elseway you could easily lost of a few fingers, maybe one eye and a lot blood. So be clever.
Converting Cs/Rb-nitrates to carbonates with carbonic acid is impossible. Where did you study chemistry?
Dense solvents are usually halogen containing e.g.: chloroform, carbon tetrachloride, trichloroethylene ect. And all of these WILL REACT with Cs, Rb
and even with Mg. But may I ask why do you need dense solvents? If the Cs floats on the surface of something than I will give you a 100% for it's
autoignition.
A little more info from Rb/Cs chemistry: mixing metallic Cs with benzene will yield a black amorphous participate what will explode on contact on
almost everything, it has a chemical formula: C6H6Cs6.
Mixing metallic Cs with toluene will also end up with a reaction and yield C6H5CH2Cs. So be smart what kind of solvent would you want to use.
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APO
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I'm talking about using carbonic acid to make carbonates from hydroxides, not nitrates, sorry for being non specific. I'm looking for a solvent that
is unreactive with caesium and rubidium so that they will coalesce easier, and I know that they won't auto ignite if I use an inert atmosphere. So if
you think that won't work how can I make make rubidium and caesium from their carbonates efficiently and safely?
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blogfast25
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| Quote: Originally posted by APO | | I would rather use carbonic acid to produce carbonates. I didn't look at the sticky thread so much because I had trouble finding dense solvents, can I
get an example of a solvent at least twice as dense as water? |
The type of inert solvents you need for reducing alkali hydroxides with Mg ALL have densities close to 1. To increase density you’d need to
introduce heavier exo-atoms into the molecules but then you can forget about inertness. Even fluorinated solvents react with hot alkali metals.
| Quote: Originally posted by APO | | I'm talking about using carbonic acid to make carbonates from hydroxides, not nitrates, sorry for being non specific. I'm looking for a solvent that
is unreactive with caesium and rubidium so that they will coalesce easier, and I know that they won't auto ignite if I use an inert atmosphere. So if
you think that won't work how can I make make rubidium and caesium from their carbonates efficiently and safely? |
Why do you want to use carbonates? The patent calls for hydroxides for a reason: that they can form thermodynamically favourable MgO in the presence
of a catalyst like t-butanol.
Most fully saturated hydrocarbon solvents are totally unreactive towards alkali metals.
[Edited on 31-12-2012 by blogfast25]
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APO
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Maybe if I do a reduction of caesium/rubidium hydroxide with magnesium, then I could take the smaller spheres and coalesce them later. I wanted to use
carbonates because they aren't very strong bases, I'm told molten caesium hydroxide
will even eat through boroscilicate glass.
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blogfast25
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| Quote: Originally posted by APO | | I wanted to use carbonates because they aren't very strong bases, I'm told molten caesium hydroxide will even eat through boroscilicate glass.
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Molten CsOH maybe, but at 200 C you're still a long way off the MP. Some superficial damage to your glass may happen though: one or
two experimenters reported that with KOH.
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Valentine
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If you have Cs2CO3 or CsOH, you can turn it into CsN3 by hydrazoic acid. The caesium azide will decompose on heating
(under N2 or argon atmosphere) to Cs and N2. I have a good method for preparing hydrazoic acid, I have attached it
(Inorganic Syntheses 1939, 1, 78-79.) If you have some ethereal solution of HN3, just pour it onto some
concentrated Cs2CO3 or CsOH solution, then decant the ether (so you can recycle it) and evaporate the CsN3 solution.
Of course, the aqueous method could be more efficient, but I think it's a bit more dangerous, because for CsN3 you don't have to distill or
dry the ethereal HN3, but you have to distill the aqueous solution and I don't like the idea of hot HN3... I found some
literature about the decomposition of the alkali azides, I have attached it, too (Zeitschrift für anorganische und allgemeine chemie 1926,
152, 52-58.)
Finally, I have found some paragraphs in the Kirk-Othmer Encyclopedia of Chemical Technology about rubidium and cesium, the latter mentions
the azide way of obtaining pure Cs at 390 °C.
I have a few kgs of caesium chloride, other cesium salts, and some sodium azide too, so I will try this method soon. I think it's a way more simple
and clean method than the reduction of molten CsOH or Cs2CO3 under inert atmosphere, if you have NaN3 and you don't
get carried away with the amount of CsN3 to decompose at once.
It is my first post at ScienceMadness, I hope you find it useful.
Attachment: Hydrazoic acid.pdf (86kB)
This file has been downloaded 1595 times
Attachment: Alkali azide decomposition.pdf (419kB)
This file has been downloaded 658 times
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APO
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Thanks!
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Valentine
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Some cesium azide is ready 
I found a good article about the decomposition process.
Attachment: Cesium azide decomposition.pdf (575kB)
This file has been downloaded 2429 times
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kristofvagyok
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This recipe is also in Georg Brauer-s book: "Handbook
of Preparative Inorganic Chemistry". -according to my memory but I never ever tried it because of the potential explosion.
If you have made it successfully please post the results, I also want to try it out(:
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Valentine
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I could do some preliminary experiments with cesium azide today. I used 2 g azide in a test tube with ground glass joint, a bath temperature of 400
°C, N2 atmosphere, a little piece of pure copper as catalyst, 10 mbar vacuum and I listened to classical music the whole time. Bubbling began just
after the salt had melted. Slowly the colorless liquid turned to dark blue. After half an hour the bluish coloration started to fade, and the liquid
became golden yellow. The cesium creeps to the walls of the tube pretty high, and it refuses to flow down. Next time I will use a lot more suitable
apparatus, and I will distill the cesium directly into an ampoule.
Azide in the bath
Temperature
Bubbling
The result
The result
I forgot taking pictures when the liquid was blue, so I will take some next time.
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elementcollector1
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Is there any condensed cesium in there, or has it all mirrored onto the tube walls?
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APO
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Can someone explain hydrazoic acid procution to me? The pdfs are sorta confusing.
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Valentine
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No, it was just a really thick mirror on the whole tube, next time I will use 15-20 g azide, and a better apparatus to get an ampoule of cesium. I
still think as a preliminary experiment it was quite good.
You put some concentrated aqueous solution of NaN3 into a RBF, then add some ether and you will get a two phase system. The HN3
is a weak acid, so if you add sulfuric acid to the system, a lot of HN3 will form in your lower (aqueous) layer. HN3 is quite
apolar, so most of it will go into the upper (ethereal) layer. If you add enough H2SO4, you will get an almost pure aqueous
Na2SO4-H2SO4 solution as a lower layer, and an ethereal HN3 solution as an upper layer. If you
have a solution of a strong base (for example CsOH), just pour the upper layer from the previous flask onto it, and as a weak acid, the HN3
will neutralize the strong base (CsOH) and goes into the lower layer forming a solution of an azide salt (CsN3). It's just playing with
polarity and acid-base equilibria.
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APO
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If I mix sodium bisulfate and sodium azide and heat it to the melting point of the sodium bisulfate, will I be able to distill some hydrazoic acid
over?
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kristofvagyok
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Valentine, this is awesome! Congrats for the Cs!
What do you think about using a paraffin wax what would only melt under 400C? I would use it as a reaction medium and a protecting from the
evaporation of the Cs.
| Quote: Originally posted by APO | | If I mix sodium bisulfate and sodium azide and heat it to the melting point of the sodium bisulfate, will I be able to distill some hydrazoic acid
over? |
Everything you learned about chemistry in school was either a) wrong, b) a crude approximation, because hydrazoic acid will explode under those
conditions.
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-Pictures from chemistry, check it out(:
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Valentine
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At 400°C cesium might react slowly with the forming decomposition products of wax. If the wax is stable enough, the cesium might endure the entire
procedure, but you will get a messed up tube in either case. How would you dispose of the wax after the decomposition?
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elementcollector1
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This might be the way to go for most home chemists, but... Cesium and oxygen... plus the availability of hydrazoic acid and/or ether, and cesium salts
in the first place... I'm very conflicted as to trying this out. Also, it's very wet where I live, if you know what I mean. I would imagine the way to
go would be to distill cesium under vacuum into a pre-formed ampoule, and seal it up while still under vacuum for best results. Or you could do what
Zan Divine did and distill from a mixture of Li and CsOH under vacuum/argon in a steel distilling apparatus leading into a RBF, and then, while still
under vacuum/argon, somehow get your sample into a good-sized ampoule. I don't know the specifics, and Zan's been out for a while (contacted him
recently, he's not dead), so... Any further ideas?
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APO
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What's the safest way to make HN3?
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Valentine
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The NaN3 - H2SO4 ehtereal method without doubt. In a previous post I described it quite well.
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APO
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Any way I can make HN3 without H2SO4?
I'm so damn scared of sulphuric acid.
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Valentine
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Don't be scared, sulfuric acid is your friend
If you still in fear, you could use 85% phosphoric acid
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ScienceSquirrel
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If you are scared of sulphuric acid you should not attempt to make hydrazoic acid.
Although it is a weak acid it is quite toxic and can explode.
http://en.wikipedia.org/wiki/Hydrazoic_acid
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APO
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I'm scared of sulphuric acid because of it's corrosiveness and that it's a liquid,so unlike most bases I can just scoop it back up, but with sulphuric
acid it would probably do some damage.
However I'm experienced with toxic, explosive, and mildly corrosive compounds.
Also under what conditions will hydrazoic acid explode?
[Edited on 23-1-2013 by APO]
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APO
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Also I decided that instead of using HN3 to make caesium azide to make caesium, that I could react magnesium and caesium carbonate under a vaccum at
200 celcius and then distill the caesium at 482 fahrenheit.
However to make my rubidium I may use the azide method.
[Edited on 23-1-2013 by APO]
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EdMeese
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Anybody have any experience with Na-B''ASE solid electrolyte systems? 99.999% pure sodium from a ZnCl2-CaCl2-NaCl bath in one report, at 200 °C so
lower energy Green blah blah. I've been obsessed for some time with the idea of doing this on an NaOH or NaOAc system closer to 100 °C -- still want
to be able to siphon off pure liquid sodium for continuous operation -- and ready to take the plunge to build one. There's an outfit in the UK that
sells BASE ceramic shapes but they start at ~100$ so I thought I'd try to make one with this Indian microwave method: Solid State Ionics 158 (2003)
199– 204. Thoughts?
"The conventional processing of sodium beta alumina (SBA) which includes synthesis and sintering in order to obtain dense bodies for use as a solid
electrolyte usually requires high temperatures (f1873 K) and long holding times. A substantial amount of Na2O may be lost during such high temperature
heat treatment, the consequence of which may be deleterious when SBA is used for device applications. To overcome this problem, the use of microwave
as heating source for the processing of SBA was attempted in the present process. Preliminary investigations on commercial Na-h-alumina bearing the
composition NaAl11O17 were carried out and it was found that Na-h-alumina was a very good absorber of microwaves. Following this, a microwave active
precursor of MgO-stabilized Na-hW-alumina was synthesized through a novel solution technique, which on compaction and exposure to microwaves yielded
dense (94% theoretical) and phase-pure Na-hW-alumina body. This paper presents the results of a direct single-step synthesis cum sintering of
MgO-stabilized Na-hW-alumina."
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