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Author: Subject: Cesium from CsCl
h0lx
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[*] posted on 4-11-2006 at 07:06
Cesium from CsCl


I could swear in my mothers name, I posted it before, but the thread vanished? Or was it deleted? Please inform me by U2U if you are going to delete this, so I won't post the third time.

I am about to obtain 100g of Cesium Chloride and I am willing to melt and electrolyse it. But I need some kind of tips or proceedures, I don't have the possibility to experiment much, because the salt is really expensive. I was thinking about some bell type thing at the cathode, with either vacuum or inert atmosphere. All tips/ideas/proceedures/drawings/picture/anything on the subject appreciated.
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[*] posted on 4-11-2006 at 07:20


Try it at first with NaCl or NaOH, and only after you can succesfully extract sodium, you may consider extracting cesium as it is much harder to do.

[Edited on 4-11-2006 by chromium]

[Edited on 4-11-2006 by chromium]




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[*] posted on 4-11-2006 at 09:35


CsCl melts at 646 C, Cs metal boils at 669C. Without propper temp control you are going to boil away the cesium.

If I was in the market for making cesium, I would go the way industry does it, distill a cesium salt with sodium.

I don't need to tell you that your inert atmosphere must be perfect, cesium catches fire in air readily.




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[*] posted on 4-11-2006 at 13:18


I am reminded of US4725311. The relevant part is quoted below.

Quote:
EXAMPLE 5

Production of Cesium Metal From Cesium Hydroxide

In the procedure described in FIG. 1, a mixture of 32.5 g cesium hydroxide (217 millimoles), 5.4 g magnesium chips (220 millimoles) and 300 ml undecane are heated with stirring at the boiling temperature of 196.degree. C. in the procedure of Example 1. 1.6 g (22 millimoles) tertiary butanol dissolved in undecane are then added in drops during 25 min and the contents of the reactor is heated with refluxing for additional 13 hours (H.sub.2 O evolution: 22 millimoles). After cooling to temperatures below the melting point of cesium, the reaction product mixture is poured onto a fine-mesh sieve to separate cesium regulus having a silvery luster from the magnesia. When the molten cesium has been filtered through a frit, the pure metal is obtained in a yield of 11.2 g, corresponding to 39% of theory.

EXAMPLE 6

Production of Cesium Metal From Cesium Alkoxide

In a reactor as described in Example 1, 69.6 g cesium hydroxide (443 millimoles CsOH, 176 millimoles H.sub.2 O) and 18.0 g (741 millimoles) magnesium chips in 300 ml dodecane are heated. with stirring at the boiling temperature (216.degree. C.) for 1.5 hours. The H.sub.2 evolution resulting from the dehydration amounted to 180 millimoles.

Thereafter, a mixture of 33.0 g (445 millimoles) tertiary butanol and 50 ml dodecane is added in drops within 2 hours and the mixture is then maintained at the boiling temperature for an additional hour. This resulted in an evolution of 440 millimoles H.sub.2 owing to the formation of cesium tertiary butylate.

The reaction product mixture consisted of a grey suspension and contained cesium spheres, which has a silvery luster and had been formed by the reaction of the cesium tertiary butylate with the surplus magnesium metal (122 millimoles). When the reaction product mixture had cooled to about 20.degree. C. it was filtered through a sieve to separate the cesium metal. Yield: 11.7 g (88 millimoles)=36% of theoretical yield. The filtrate that had passed through the sieve was filtered through a frit. The residue retained on the frit was washed several times with tetrahydrofurane.

From the frit filtrate containing 276 millimoles cesium butylate and 84 millimoles magnesium butylate, pure cesium butylate could be recovered by recrystallization from a mixture of tetrahydrofurane and toluene.


If you're interested, look at the Unconventional Sodium thread. There was some discussion about this patent there.




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[*] posted on 4-11-2006 at 13:59


Alternatively, you can use the method in Brauer which is (for me) the best available. CsCl is heated with calcium powder or turnings under high vacuum in a thrice purged system that must be constructed of either stainless steel or a temperature resistant glass. This is the method I will use when/if I get around to doing it. I happen to have a friend who is supplying the cesium chloride. If you'd like, feel free to U2U me about it. The yield is 98% of theory. I should remind you that electrolyzing it is not easy since the metal is soluble in its salt. You could however electrolyze the molten salt above the B.P. of Cs and distill over cesium, but you still will have CsCl in the metal: the salt is volatile, especially in vacuum. Also, reducing with sodium has inherent problems as well if conducted in vacuo: some sodium inevitably distills over as well.


Brauer's is availabe in the library.




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[*] posted on 5-11-2006 at 02:40


How can Ca replace Cs? It is much less active than Cs.
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[*] posted on 5-11-2006 at 05:33


Cs is more volatile. An equilibrium will exist between Ca and Cs. However, Cs can be driven off as gas, and Ca remains behind. This causes the equilibrium to be driven to the Cs side:

2Cs(+) + Ca <-->>>>> 2Cs + Ca(2+)




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[*] posted on 5-11-2006 at 06:52


I have read, but after a search can not locate the reference, of the production of caesium metal from the reaction of CsF with an excess of iron powder in a stainless steel retort under reduced pressure (and after argon flush, before pumpdown)

2 CsF + Fe => FeF2 + 2 Cs (gas)

FeF2 has a fairly high boiling point, something like 1800 C, CsF is 1250 C or so and melts a little above the boiling point of Cs at 1 atmosphere Again this is a case of forcing the reaction through removal of a product.
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[*] posted on 5-11-2006 at 22:34


I'm wondering if you could do the cesium or rubidium separation/distillation in an a quartz apparatus instead of borosilicate since I hear that lithium will attack and consume borosilicate at high temps. do the other alkali metals react with borosilicate as well so that you can't use it for distillation ? would quartz work or be similarly affected?

[Edited on 6-11-2006 by Maya]
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[*] posted on 6-11-2006 at 00:10


There is also a method of heating cesium dichromate with zirconium powder in vacuum that is stated to be the most convenient means of cesium production, as the temperature only needs to be at the boiling point of Cs and not much higher (no equilibrium has to be shifted). I think this is also outlined in Brauer, look it up.



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[*] posted on 6-11-2006 at 06:50


And where could one obtain Cs2Cr2O7? The only dichromates, which are available for the home chemist (and even those with some difficulty in many countries) are the sodium, potassium and ammonium salts.

Or is there a nice synth of cesium dichromate from the other dichromates. Is Cs2Cr2O7 less soluble, so that it can be crystallized? This subject is interesting to me, because I hope to receive 100 grams of CsBr next week or so.

[Edited on 6-11-06 by woelen]




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[*] posted on 6-11-2006 at 07:00


As potassium dichromate is significantly less soluble that sodium dichromate, I would expect the cesium salt to follow the rule and be even less soluble (with perchlorates it follows this trend, KClO4, RbClO4 and CsClO4 are progressively less soluble).
I have no reliable data however.

[Edited on 6-11-2006 by garage chemist]




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[*] posted on 6-11-2006 at 07:34


It may be that garage chemists idea does not work as cesium chromate is quite well soluble in water. Do not know about dichromate though.

Some cesium salts can be made by reacting cesium sulfate with barium salt of desired anion. Cesium hydroxide can also be made this way.

Cesium sulfate can probably be prepared by heating cesium halide with sulfuric acid and lettting volatile acid to escape but it may be quite hard to go beyond hydrosulfate this way.




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[*] posted on 6-11-2006 at 16:15


Woelen, look in Brauer's, there is a section on the formation of rubidium and cesium (di)chromates. The method chromium mentioned is what is actually used. Barium dichromate is added to a warm solution of Cs2SO4 (or that of rubidium).



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[*] posted on 6-11-2006 at 16:55


Quote:
Originally posted by Maya
I'm wondering if you could do the cesium or rubidium separation/distillation in an a quartz apparatus instead of borosilicate since I hear that lithium will attack and consume borosilicate at high temps. do the other alkali metals react with borosilicate as well so that you can't use it for distillation ? would quartz work or be similarly affected?

[Edited on 6-11-2006 by Maya]


Yes. There is plenty of literature on the distillation of K/Na/Rb/Cs in borosilicate/quartz.
Lithium is fairly unique in that respect. No worries in that respect.

Does anyone want to explain to me why CESIUM is so much cheaper than RUBIDIUM? Is it just because of the relative abundances?

(Edited so it actually makes sense)

[Edited on 8-11-2006 by iamthewaffler]




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[*] posted on 8-11-2006 at 00:18


Quote:

Does anyone want to explain to me why rubidium is so much cheaper than cesium? Is it just because of the relative abundances?



Seems around here the opposite, cesium is much cheaper then rubidium
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[*] posted on 8-11-2006 at 00:25


Quote:
Originally posted by Maya
Quote:

Does anyone want to explain to me why rubidium is so much cheaper than cesium? Is it just because of the relative abundances?



Seems around here the opposite, cesium is much cheaper then rubidium


Um. That's what I meant. I feeeeeel dumb.




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[*] posted on 24-12-2010 at 18:38


Quote: Originally posted by iamthewaffler  

why CESIUM is so much cheaper than RUBIDIUM? Is it just because of the relative abundances?
Only cesium has own minerals with high Cs content and therefore easily winnable Cs (e.g. Pollucite (Cs-Na-Silicate), Pautovite (Cs-Fe-Sulphide), Cesstibtantite (Cs-Na-Sb-Ta-Oxide)), rubidium is only present in other minerals (small % concentrations in Leucite, Lepidolite, others) and has no known own mineral (can only be produced as a side product). - This is often the case for production of elements: concentration in a mineral is more important than total concentration in earth crust. (e.g. gallium is as abundant in the earth crust as lead, but many times more expensive, because lead forms many minerals = high local concentration = easily winnable....gallium forms very rare minerals, is mostly present in a scattered pattern))

So: relative abundances are not the reason: wikipedia.de says: Cs = 6.5 ppm, Rb = 29 ppm (if these values aren't correct: other sources at least say: "Rb > Cs"

Another reason might be: Cs is used in a larger amount (and thus produced) in science and technology than Rb. Probably because chem/phys. Cs-K difference is larger than Rb/K (???).
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[*] posted on 25-12-2010 at 00:26


Hm, after the sucess with "pok's" method for Potassium, maybe you guys should try with NaOH and then with CsOH, as according to the same patent it should work.
That will be a pretty big breakthrough in amateur science.
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[*] posted on 25-12-2010 at 06:43


Quote: Originally posted by Cuauhtemoc  
Hm, after the sucess with "pok's" method for Potassium, maybe you guys should try with NaOH and then with CsOH, as according to the same patent it should work.
That will be a pretty big breakthrough in amateur science.


Yes but we need first to tackle the supposedly easier sodium with the alkoxyde - Mg reduction method. So far no one here has even attempted, presumably because the patented reaction times are so high. Issues with the solubility of sodium t-butoxyde, we feel...
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[*] posted on 25-12-2010 at 06:52


Quote: Originally posted by Cuauhtemoc  
Hm, after the sucess with "pok's" method for Potassium, maybe you guys should try with NaOH and then with CsOH, as according to the same patent it should work.
That will be a pretty big breakthrough in amateur science.


Indeed, however Cs is such a scary, reactive metal, that it would have to be done in a professional laboratory under the most ideal of conditions, in a noble gas atmosphere, with the most precise of temperature controls and the highest purity of anhydrous reagents. As tempting as it would be, I would never try this in a "home lab"... too dangerous.

There are elements or chemicals that I don't mind doing without, simply because they are above my level of comfort. Cesium and Rubidium are some of them.

Robert
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[*] posted on 25-12-2010 at 09:24


Quote: Originally posted by Arthur Dent  

There are elements or chemicals that I don't mind doing without, simply because they are above my level of comfort. Cesium and Rubidium are some of them.

Robert


I take it you're OK with fluorine then! :P
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[*] posted on 25-12-2010 at 12:15


Yup!
I have two or three cups of chlorine trifluoride every morning, makes my stomach a bit upset though. ;)

Yeah, I have to admit I never take any chemicals for granted, and before I proceed with an experiment, study its mechanisms and interactions thoroughly. There are many chemicals that I deem too reactive, noxious or plain dangerous for me to have, like elemental phosphorus, cyanide compounds, volatile toxic stuff and so on. I have limited space and I can't have a cool lab with a fume hood so I do my "smelly" stuff outdoors! LOL

Plus many chemicals would serve no purpose for my basic chemistry usage, mainly electroplating, etching, precious metal recuperation and geeky, fun experiments.

Robert


Robert

[Edited on 25-12-2010 by Arthur Dent]
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[*] posted on 25-12-2010 at 14:20


In my experience, K/Na alloy is worse than elemental cesium because it is spontaneously flammable with air and can generate much more hydrogen gram for gram than cesium can. Cs and Rb might react "quicker" but they're stoichiometrically less dangerous with water. Both can be handled in a glove box (or a bag filled with Ar). Treat it with respect as you would any pyrophoric material and you'll be fine.

While KOH may not etch glass at that temperature, I'm willing to bet that CsOH will. CsOH is much more corrosive than KOH.




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[*] posted on 25-12-2010 at 18:00


Quote: Originally posted by h0lx  
I am about to obtain 100g of Cesium Chloride and I am willing to melt and electrolyse it. But I need some kind of tips or proceedures, I don't have the possibility to experiment much, because the salt is really expensive. I was thinking about some bell type thing at the cathode, with either vacuum or inert atmosphere. All tips/ideas/proceedures/drawings/picture/anything on the subject appreciated.


Tip? Sure answers to you question can be found in —

Mellor's Comprehensive Treatise on Inorganic and Theoretical
Chemistry Volume II Supplement III p. 2287 & ff.

75% yield by heating the chloride w/ calcium carbide in vacuo
@ 700-800o seems the easiest. However, you may find one of
the other methods better suited to your I got's.

The barium azide method seems upon me — the most exciting...!


Say - Whatever happened to the a in cerium?
i In aluminum? a in archeology?

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