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Author: Subject: Turning CsNO3 to other cesium salts.
Mixell
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[*] posted on 4-10-2011 at 10:26
Turning CsNO3 to other cesium salts.


I was thinking about buying 200g of CsNO3, but I actually need the chloride salt (or hydroxide/carbonate/ any other form that can be turned into other salts) of cesium.
I know that aqua-regia decomposes in the following reaction: HNO3 (aq) + 3 HCl (aq) → NOCl (g) + Cl2 (g) + 2 H2O (l)

I was thinking about adding HCl to the cesium nitrate (which I do not know if its soluble in, although CSCl is soluble in concentrated HCl) and creating a poor man's aqua-regia, but I do not know if the decomposition of the nitrate ions will occur (and how fast).
If it indeed occurs, I need top find a way to speed it up and hopefully decomposes all of the nitrate with excess HCl, any suggestions about that?
Or about other methods of forming CsCl from CsNO3 (precipitating it as CsClO3/CsClO4 is not an option, I do not posses salts with those anions at all).
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[*] posted on 4-10-2011 at 11:39


Why don't you buy the chloride? AFAIK, it's much easier to get and should be cheaper. Caesium nitrate is not so common, and the chloride is used in biochemical laboratories a lot.



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[*] posted on 4-10-2011 at 11:52


"As with other alkali metal nitrates, caesium nitrate decomposes on gentle heating to give caesium nitrite:
2CsNO3 → 2CsNO2 + O2"
from
http://en.wikipedia.org/wiki/Caesium_nitrate
which suggests that, on strong heating, you will get Cs2O (as long as you have a container that will stand the hot oxidising base)

You can then convert that to any salt by adding water (cautiously) to get CsOH then the appropriate acid.

Or, you can decompose the nitrite by adding HCl The advantage would be that you need less HCl.
In any event, be aware that the Cl2/ NOx/ NOCl2 gases produced are toxic and corrosive.
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[*] posted on 4-10-2011 at 13:08


CsNO3 + HCl + SbCl3 === > CsSbCl4 + HNO3. The CsSbCl4 is poorly soluble, decant off the HNO3 and wash with cold water.

Treat slurry of CsSbCl4 with H2S:

CsSbCl4 + 3/2 H2S === > CsCl + ½ Sb2S3 + 3 HCl. Filter to remove insoluble Sb2S3 and evaporate to get pure CsCl. Luvvly jubbly with all that H2S needed! (but it works: classical work up of Cs salts from minerals).




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[*] posted on 4-10-2011 at 13:35


I prefer not to mess with H2S... :)
May be the same thing can be done with copper or similar metals? I actually need the CsCl to create some salts, like CsCuCl3 or Cs2CuCl4, although I don't know if the nitrate ion will interfere in the reaction.
Also, I can't achieve temperatures that can decompose CsNO2 to Cs2O (not even close to that).
CsCl cost quite a lot (as CsNO3), but I found 200g of CsNO3 for 30$ including shipping, I think its quite cheap...
So you say that its better to decompose the nitrite (although the nitrate decomposes at least at 414C, and that's also quite high...) with HCl instead of decomposing the nitrate using the same method?
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[*] posted on 4-10-2011 at 18:00


I think you're on the right track with common ion'ing it to death. Dissolve as much as you can in aqueous HCl and then gas the heck out of it with HCl, similar to the purification of NaCl. Anything that precipitates should be the chloride. Evaporate off some HCl laden solution, gas again and collect. Continue until you run out of liquid and start over. HNO3 should be removed as well by the procedure both from the evap and the slow decomp.

There is also a similar method (only in method of execution, not chemistry) to make perchloric acid. Ammonium perchlorate is heated with nitric acid to decompose the ammonia part of the molecule and leave behind the perchloric acid. This is usually done by heating the solution far from civilization and adding in nitric acid as the solution volume dwindles. In this way you constantly supply the nitric the reaction needs as well as running at high temps to ensure fast reaction. If you dissolved your cesium nitrate in hydrochloric acid and then heated and continued to evaporate, replenishing with HCl as the liquid level dropped, it won't matter that the HCl azeotrope should come off first, you will continually remove the nitric acid and eventually take it completely to the chloride salt. The only challenge being to secure a relatively pure supply of hydrochloric acid free of metals.

A little more on the wild side might be to treat with hydrochloric acid then try to extract the HNO3 with DCM. Although I know HNO3 has been extracted into DCM I am not sure this can be done with dilute solutions. Just spit balling here.





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[*] posted on 4-10-2011 at 23:35


Could CsN3 be made from the oxide or nitrate salt using ammonia?
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[*] posted on 5-10-2011 at 01:39


Quote: Originally posted by MyNameIsUnnecessarilyLong  
Could CsN3 be made from the oxide or nitrate salt using ammonia?


Nop. But I think you meant Cs3N, right? Akaline metals do not form stable nitride. Well, lithium does, but I don't think that caesium would. Maybe...
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[*] posted on 5-10-2011 at 02:03


No, I was referring to the azide. I'm trying to get cesium metal, and decomposing the azide salt seems to be the easiest method for making elemental cesium. But it's hard to find it cheaper than just buying the metal itself.
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[*] posted on 5-10-2011 at 08:08


I will try doing it with NaNO3 first. Decomposing the NaNO3 with HCl, and decomposing NaNO3 to NaNO2 via heating and decomposing the NaNO2 with HCl too.
I also thought about using H2SO4, getting the sulfate salts, and then using BaCl2 that will form a precipitate of BaSO4 and will leave a solution of CsCl.
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[*] posted on 5-10-2011 at 10:02


To reduce the amount of conc. HCl required for distilling off the HNO3, you could first decompose the nitrate by refluxing a mixture of CsNO3, 20% HCl(aq) and some ethanol until no more nitrogen oxides fumes form and then remove the HCl(aq) excess and the volatile organic side products. Another portion of conc. HCl and evaporation should make sure to remove all the nitrate/nitrite residues. Already HCl is able to reduce nitric acid in acidic enough media (e.g., aqua regia), but adding ethanol would make sure no nitrate remains after solvent removal.

Of course, making CsBr and CsI from CsNO3 by using HBr(aq) or HI(aq) correspondingly should be much easier as the bromide and iodide both easily reduce nitric acid to yield volatile products.




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[*] posted on 5-10-2011 at 10:39


So the ethanol will acts as a reducing agent? And the HCl as a proton source.
But the problem is that I don't have ethanol.
I have other volatile organics though: acetone, formic acid, acetic acid and hexane. I know formic acid is a reducing agent, although I'm not sure about the rest, can anyone of them be used instead of ethanol?
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[*] posted on 5-10-2011 at 11:30


I once did some work looking into an explosion that happened when someone mixed formic and nitric acids.
Do you live somewhere that doesn't allow alcohol? Vodka is good enough for this sort of thing.
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[*] posted on 5-10-2011 at 11:50


Vodka is for drinking! :)
I do have some cheap vodka around somewhere, but I don't know if it will be enough, and I hope it doesn't have any impurities in it.
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[*] posted on 5-10-2011 at 12:41


The cheaper, the better in regard to the vodka used for such a purpose.
Formic acid will also do provided you use a more diluted HCl(aq) like 2-4M or thereabout. I would expect it to get oxidized much more rapidly than ethanol, hence the accident unionised talks about (I wonder why someone mixed concentrated formic and nitric acids?). However, in a diluted mixture, the reaction rate and energy released should be tolerable (beware of spillovers due to large amounts of CO2 bubbling). Try with 3-6 equivalents of ethanol or formic acid (I would still rather opt for ethanol/vodka). It should be enough. I think a couple of hours of reflux should be OK. You should be able to see some brown fumes of nitrogen oxides leaving the condenser and getting oxidized by air to the brown NO2.
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[*] posted on 6-10-2011 at 02:31


Denaturated alcohol ('methylated spirits') is available from just about any hardware store, probably cheaper than Vodders...



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[*] posted on 7-10-2011 at 12:13


If you have means to create 1000C temperatures, you could try a modification of the <a href="http://en.wikipedia.org/wiki/Leblanc_process">Leblanc process</a>. (I am assuming that caesium chemistry is sufficiently similar to sodium and potassium chemistry for this to work.)


  • Convert the caesium nitrate into caesium sulphate using sulphuric acid. (This type of reaction is discussed abundantly in the nitric acid threads.)
  • Then heat the caesium sulphate with calcium carbonate and charcoal at 1000C, to give caesium carbonate, calcium sulphide and carbon dioxide.
  • Extract the caesium carbonate, by dissolving in water, and filtering everything else out.
  • React with hydrochloric acid.




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[*] posted on 8-10-2011 at 01:19


I would not use anything else than HCl. Ethanol leads to formation of acetate, which is hard to remove, formic acid may lead to formation of oxalate and remains of formiate. HCl is cheap and easy to obtain. Just use that. I would first distill the HCl if it is hardware store grade, otherwise you get all kinds of impurities in your CsCl which may be hard to remove.



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[*] posted on 8-10-2011 at 12:23


Woelen:

Hmmm. Nitrate ions with strong HCl and some ethanol/methanol should give ethanoic/methanoic acid mixture. On boiling to dryness this should evaporate, leaving behind pure chloride. Bit of a waste of nitrate and good alkanol but there your go!




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[*] posted on 9-10-2011 at 03:46


I've tried several reduction/decomposition methods with NaNO3 and HCl (all in room temperature):
Acetone does nothing at all.
Formic acid produces a bit of bubbling, might be good during a period of several daysת EDIT: On heating it works very well too.
Poor man's aqua regia also produces a tiny amount of bubbles, could be practical at elevated temperatures.
And finally, ethanol produces a self sustainable reaction, but the test tube is quite hot, so I got a feeling quite a bit of HCl/ethanol is evaporated.
My conclusion is: Vodka, you did it again!

But I'm confused as to how much ethanol should I use for my large scale experiment, I know each mole of ethanol give 4 moles of electrons upon oxidation to acetic acid, but can't acetic acid be further oxidized to CO2 and H2O?

EDIT: I'm still not sure if I should use formic acid or ethanol.
The formic acid seem to work even better.


[Edited on 9-10-2011 by Mixell]

[Edited on 9-10-2011 by Mixell]
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[*] posted on 9-10-2011 at 04:41


Well, formic acid seems like the most reliable option, it works fast, leaves behind CO2 and water and its evaporation temperatures is around 100C. The only disadvantages is that the reaction is very vigorous, a few ml's produces a lot of heat and vigorous bubbling, so the large scale experiment should be conducted via small additions of formic acid and/or HCl.
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[*] posted on 9-10-2011 at 05:11


Did you test the dry remains for purity? Try dissolving some of this in water and add a solution of CaCl2. It should remain clear. If not, then there may be some oxalate left over in the residue.

@blogfast25: I do not fully agree with what you say. I once tried a similar reaction of ethanol with dichromate and some acid in order to get the corresponding chromium salt of the acid, mixed with the potassium salt of that acid. I just let the mix react until I had a pure green color and then evaporated to dryness (in this case, I obtained a kind of green syrup which solidified to a glass-like solid). Lateron, I found that my mix also contained some acetate. Not much, but enough to notice its presence (faint smell of vinegar on addition of a strong acid). So, I can imagine that with cesium you also get some acetate.




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[*] posted on 9-10-2011 at 06:29


I did not test it, because I didn't used stoichiometric amounts. On the larger scale I will try to use it.
May be it is possible to use H2O2 to oxidize any left over oxalic/formic acid to CO2 and water (in presence of concentrated HCl of course) and just evaporate everything afterwards.
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[*] posted on 9-10-2011 at 09:49


Quote: Originally posted by woelen  
@blogfast25: I do not fully agree with what you say. I once tried a similar reaction of ethanol with dichromate and some acid in order to get the corresponding chromium salt of the acid, mixed with the potassium salt of that acid. I just let the mix react until I had a pure green color and then evaporated to dryness (in this case, I obtained a kind of green syrup which solidified to a glass-like solid). Lateron, I found that my mix also contained some acetate. Not much, but enough to notice its presence (faint smell of vinegar on addition of a strong acid). So, I can imagine that with cesium you also get some acetate.


Similar to my experience with K2Cr2O7 and ethanol (and sulphuric acid): my KCr alum didn't want to crystalise (remember the exchange?). But I cannot see how residual acatate can stay in solution: there's no cations to account for it... In the case of this Cs displacement, since as acetates are so soluble, a recrystallisation (recommendd anyway) of the CsCl dhould eliminate any doubts.

[Edited on 9-10-2011 by blogfast25]




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[*] posted on 9-10-2011 at 11:24


Yes, with the latter you are right. If any acetates remain in the CsCl, then a recrystallization will remove them and when a recrystallization is applied then the use of ethanol to destroy the nitrate/nitric acid is a faster method of destruction than using HCl alone.

Recrystallization of CsCl must be done very carefully, its solubility is amazing. CsCl, however, is only somewhat hygroscopic, so doing this on a warm dry place such as a heat radiator certainly should be possible. The residue, containing remaining Cs can easily be recovered by adding a solution of a perchlorate, or if that is not available, a chlorate or bromate.




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