CaCl2
Harmless
Posts: 39
Registered: 14-1-2017
Member Is Offline
Mood: No Mood
|
|
CsCl into CsOH or Cs2CO3?
Cesium chloride is probably the most available cesium salt, not exactly common, and is pretty expensive, but you can still find it on ebay.
It is also pretty hard to use as a strating point for making other cesium salts, if you could get turn it into the carbonate or hydroxide you could
then easily make other salts just by adding an acid.
Some of the things I have seen being suggested are:
1. Add saturated KOH or NaOH solution to the CsCl solution, the sodium or potassium chloride will precipitate. This still leaves a lot of the cloride
into the solution.
2. The wiki suggests using electrolysis to oxidize the chloride into chlorine. This sounds slow, energy expensive, probably will leave a lot of
chloride into the solution. It also produces toxic chloride.
Neither of these options seems to really work, so I have been looking for a better solution.
One way of getting rid of the chloride is mixing solutions of CsCl and aluminum sulfate to precipitate cesium alum, which has a low solubility in cold
water. But I'm not sure if the alum is any easier to work with than the chloride.
Are there any good methods for making CsOH or Cs2(CO3) from CsCl?
Cesium is an uncommon enough topic that I thought this would be a better place to post this than beginings, sorry if this was wrong. I tried
searching, but couldn't find any subjects about the topic, sorry if I missed some. I'm sorry for any grammar/spelling mistakes, I'm not a native
English speaker, also sorry apologizing way too much. (No, I'm not a Canadian.)
[Edited on 29-10-2017 by CaCl2]
|
|
|
clearly_not_atara
International Hazard
Posts: 2787
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
The simplest way I think would be to make cesium phosphate and react this with calcium oxide. This requires heating CsCl with phosphoric acid until
hydrogen chloride is released.
Cesium can also be precipitated as the bitartrate IIRC if you have tartaric acid handy; tartaric acid is sold for making wine and cheese and can be
purchased on Amazon. Add tartaric acid to CsCl and raise the pH to 2-3ish by adding a solution of NaHCO3.
Cesium bitartrate can then be converted to the hydroxide by the action of calcium oxide, precipitating calcium tartrate.
If you don't mind the noxious gas you can also precipitate cesium as the nitrate and heat this to decomposition.
[Edited on 29-10-2017 by clearly_not_atara]
|
|
|
CaCl2
Harmless
Posts: 39
Registered: 14-1-2017
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by clearly_not_atara  |
Cesium can also be precipitated as the bitartrate IIRC if you have tartaric acid handy; tartaric acid is sold for making wine and cheese and can be
purchased on Amazon. Add tartaric acid to CsCl and raise the pH to 2-3ish by adding a solution of NaHCO3.
Cesium bitartrate can then be converted to the hydroxide by the action of calcium oxide, precipitating calcium tartrate.
|
This method seems most promising since I think I have all the chemicals needed already.
Do you have any info on the solubility of Cesium bitartarate? I couldn't find any with a quick search. The amount of wasted cesium depends on it, and
since cesium salts are expensive it's pretty important.
How would one go about reacting the cesium tartarate with the calcium oxide. Does it work by just adding solid CaO to the tartarate solution or does
it need to be dissolved as Ca(OH)2 first?
[Edited on 29-10-2017 by CaCl2]
[Edited on 30-10-2017 by CaCl2]
|
|
|
clearly_not_atara
International Hazard
Posts: 2787
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
Apparently caesium bitartrate has low solubility in water but not as low as the corresponding bitartrates of potassium and rubidium:
"A better method of separation is by means of the hydrogen tartrates, that of rubidium being soluble in 8'5 parts of boiling water and in 84*57
parts at 25 ; whilst caesium hydrogen tartrate dissolves in 1'02 part of boiling water and in 10'32 parts at 25 (Allen), and potassium hydrogen
tartrate requires for solution 16 parts of boiling water and 77 parts at 25."
-- from "Introduction to qualitative chemical analysis", http://anonym.to/http://www.archive.org/stream/introductiont...
Another insoluble caesium salt is sodium dicaesium hexanitritocobaltate, which may be useful if you really don't want to lose any caesium, Dicaesium
hexachlorostannate is also insoluble and results on treating a solution of Cs+ with HCl and SnCl4.
EDIT: I'm not sure how to react the bitartrate with calcium oxide, but I suggest powdering the solids if possible (or at least breaking them into
small pieces) and performing the reaction with cold water, possibly under grinding. The reagents will not go completely into solution, but the
reaction can still proceed.
[Edited on 30-10-2017 by clearly_not_atara]
|
|
|
CaCl2
Harmless
Posts: 39
Registered: 14-1-2017
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by clearly_not_atara |
"A better method of separation is by means of the hydrogen tartrates, that of rubidium being soluble in 8'5 parts of boiling water and in 84*57
parts at 25 ; whilst caesium hydrogen tartrate dissolves in 1'02 part of boiling water and in 10'32 parts at 25 (Allen), and potassium hydrogen
tartrate requires for solution 16 parts of boiling water and 77 parts at 25."
|
Sodium hydrogen tartarate apparently has solubility 89g/L at 20°C
https://www.chemexper.com/cheminfo/servlet/org.dbcreator.Mai...
That's pretty close to that of cesium bitartarate, but one presumably doesn't have to worry about precipitating it instead since the amount of sodium
would be smaller than that of cesium.
Though the whole thing is made more complicated by the various isomers of tartaric acid, no idea if the two sources are talking about the same ones.
[Edited on 31-10-2017 by CaCl2]
Edit by moderator : Fixed quote
[Edited on 10-31-2017 by gdflp]
[Edited on 1-11-2017 by CaCl2]
|
|
|
clearly_not_atara
International Hazard
Posts: 2787
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
Hm, yes that would still work but it might be difficult... I think that the D/L isomers tend to be slightly less soluble than the meso isomer but only
slightly. Most commercially available tartaric acid is the natural form IIRC.
There are a couple of improvements I can think of - either:
* use an excess of tartaric acid so you can dilute the solution of HCl somewhat without adding any other metals... you can go as high as 1:1 tartaric
acid:water
* form sodium bitartrate beforehand and use this to perform a more precise double displacement with CsCl by e.g. adding solid CsCl to a saturated
solution of NaH(CO2CHOH)2
It kind of depends on how much you're paying for tartaric acid vs. how much extra work you feel like doing. Since the solubility of sodium bitartrate
is quite low it should be possible to precipitate some from solution pretty easily... just mix one part NaHCO3 and one part (HO2CCHOH)2.
|
|
|
unionised
International Hazard
Posts: 5126
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
PPT the bitartrate.
I think you can do that using ammonium bitatrate, which would avoid adding any other metals.
Burn it to carbonate.
(any Ammonium would be lost too)
|
|
|
symboom
International Hazard
Posts: 1143
Registered: 11-11-2010
Location: Wrongplanet
Member Is Offline
Mood: Doing science while it is still legal since 2010
|
|
cesium chloride and sulfuric acid forming cesium sulfate and hydrochloric acid. calcium hydroxide is reacted with cesium sulfate to form calcium
sulfate insoluble and cesium hydroxide
This is the most otc route excluding cesium chloride that is bought online.
Also DO NOT use a glass or aluminum container
It will dissolve it
The ammonium bitatrate would be the cleanest if you can find tartaric acid
Precipitating ceasium tartrate
|
|
|
clearly_not_atara
International Hazard
Posts: 2787
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
Ammonium bitartrate is even more insoluble than cesium bitartrate:
https://en.wikipedia.org/wiki/Solubility_table
(1.88% at 10 C and 2.7% at 20 C)
In general bitartrate salts of monovalent ions are very insoluble, although sodium is (usually!) the exception. I believe lithium bitartrate also has
high solubility. Zinc forms double salts; magnesium inhibits the deposition of tartrates. Quite simply, the ion is a jerk.
Hexamethonium bitartrate on the other hand has a very high solubility because it is a double quaternary ammonium salt which cannot form a tightly
bonded crystal structure with bitartrate ions.
It's possible that lysine or arginine could be used as the base, though, since their bitartrates should be quite soluble. It should be possible to
obtain these. But I still think sodium bitartrate can work pretty well as long as you're careful with the stoichiometry.
Calcium nitrate might also work, since cesium nitrate has roughly the same solubility as cesium bitartrate, whereas Ca(NO3)2 is extremely soluble.
[Edited on 1-11-2017 by clearly_not_atara]
|
|
|
18thTimeLucky
Hazard to Self
Posts: 51
Registered: 19-8-2017
Location: The one-and-only tea and crumpet land (UK)
Member Is Offline
Mood: 0 Kelvin and still won't crystallise from solution
|
|
| Quote: Originally posted by CaCl2 |
2. The wiki suggests using electrolysis to oxidize the chloride into chlorine. This sounds slow, energy expensive, probably will leave a lot of
chloride into the solution. It also produces toxic chloride.
|
I have had a decent amount of success from this route with other alkali metal chlorides such as sodium and potassium. I cannot deny it is slow, but
there should be minimal chloride contamination in the hydroxide solution as it forms.
Most sources will say electrolysis of an alkali metal chloride solution will give you the corresponding alkali metal hydroxide, which is true, but
they always seem to present the cell as a single electrolysis cell. This will not produce hydroxides but instead hypochlorites and chlorates.
Separating the solutions though solves this problem - I did this with a tube connecting the otherwise separated two cells, and using cotton inside the
tube to act as an ion-permeable membrane. Yes, I know, cotton does funnily work (hence the yellow in the picture is only in one cell and not spreading
into the other). Separating the cells also keeps the chloride solution and the forming hydroxide solution separate.
I talk more about it on my blog (there is a link in my signature) if you are interested. For example, I ended up making 205ml of a ~1M solution of
sodium hydroxide from table salt in 1-2 days through membrane electrolysis. The whole procedure is very OTC.
If you want speed then maybe using tartrates or something else is a better idea though.
EDIT: If you don't want to deal with toxic chlorine gas, especially if you speed up the rate of electrolysis to make it more viable, a chlorine
scrubber is easily made. Just bubble the gas into a sodium hydroxide solution and convert it to sodium hypochlorite, for example.
[Edited on 1-11-2017 by 18thTimeLucky?]
Yep, I have a chemistry blog!
18thtimelucky.wordpress.com
"Amateur chemistry does seem like being in a relationship with someone very beautiful and seductive but has expensive taste, farts a lot and doesn't
clean up after themselves, but you love them anyway" - a dear friend |
|
|
clearly_not_atara
International Hazard
Posts: 2787
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
This thread may be helpful:
https://www.sciencemadness.org/whisper/viewthread.php?tid=69...
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
