Sciencemadness Discussion Board - Making other cesium salts from cesium alum

Making other cesium salts from cesium alum

woelen - 20-8-2009 at 11:24

I recently purchased 175 grams of cesium alum (CsAl(SO4)2.12H2O) for just GBP 8. I intended to use this to make other cesium salts by double displacement (e.g. cesium perchlorate, cesium bromate). But to my disappointment this salt hardly is soluble.

0.29 grams per 100 ml at 10 C
1.24 grams per 100 ml at 50 C
5.29 grams per 100 ml at 80 C

This is MUCH less than the solubility of plain potassium alum.

I did not even manage to make CsClO4, which is one of the least soluble perchlorates

Does anyone of you have an idea how the cesium can be released from this compound, such that it can be in solution and can be used to make other salts? I myself tried adding a little amount of sulphuric acid to the cesium alum, hoping that it would dissolve a little bit better (hoping that the sulfate ions would forms bisulfate and that this leads to better solubility). Unfortunately this did not work.

Phosphor-ing - 20-8-2009 at 11:35

This might help

http://www.patentstorm.us/patents/4466950/claims.html

Mentions using a hydroxide solution for cesium alum.

[Edited on 20-8-2009 by Phosphor-ing]

blogfast25 - 20-8-2009 at 12:17

Quote: Originally posted by woelen

I recently purchased 175 grams of cesium alum (CsAl(SO4)2.12H2O) for just GBP 8. I intended to use this to make other cesium salts by double displacement (e.g. cesium perchlorate, cesium bromate). But to my disappointment this salt hardly is soluble.

0.29 grams per 100 ml at 10 C
1.24 grams per 100 ml at 50 C
5.29 grams per 100 ml at 80 C

Nice purchase! Where did you buy it?

Are the solubility data yours?

blogfast25 - 20-8-2009 at 12:33

Woelen:

From this source:

http://docs.google.com/gview?a=v&q=cache:qLhS5BV_GqoJ:pu...

"The alum is roasted with 4 % carbon and leached to yield a Cs2SO4 solution which can then be converted to CsCl". (page 4: Production technologies [of Cesium])

Looks like you might have to go non-aqueous first...

12AX7 - 20-8-2009 at 14:20

It's not enough to add base and leach out Cs2SO4 (erm, and Na2SO4) from the goo?

Tim

garage chemist - 20-8-2009 at 14:41

Try adding a suitable (not too strong) base to a hot solution of the alum to precipitate the aluminium as Al(OH)3. Ammonia and sodium carbonate should both work. You need to get the pH value into the region where solid Al(OH)3 exists- it's amphoteric, but when the pH is right, it will precipitate quantitatively.
Use only a stochiometric amount of sodium carbonate (so that it gets protonated to bicarbonate).

I have a german procedure for the preparation of alumina hydrate from potassium alum that uses sodium carbonate to precipitate Al(OH)3 from an alum solution. The mixed solution is kept hot for a few hours, presumably to make the Al(OH)3 more filterable.
After thorough washing, the precipitate is dissolved in HCl and re-precipitated by pouring into NH3.

[Edited on 20-8-2009 by garage chemist]

UnintentionalChaos - 20-8-2009 at 15:07

If you were to mix a hot solution of the alum with BaCl2 and Ba(OH)2 in the right proportions, after filtration, you should be left with an essentially neutral solution of CsCl

2CsAl(SO4)2*12H2O + 3Ba(OH)2 + BaCl2 -> 2CsCl + 4BaSO4 + 2Al(OH)3 + 12H2O

The only issues I see are loss of solution in the Al(OH)3 goo if you only have gravity filtering and the dilution of the resulting solution.

Sr(OH)2 is reasonably soluble in boiling water and very easily prepared from other strontium salts and NaOH if you are looking to dodge the barium.

100% yield corresponds to 51.85g of CsCl.

len1 - 20-8-2009 at 22:53

Is this a practically checked procedure or just a ballanced formula - in the latter case then so is this

2CsAl(SO4)2*12H2O + 3Ba(OH)2 -> Cs2SO4 + 3BaSO4 + 2Al(OH)3 + 12H2O

not_important - 20-8-2009 at 23:44

Most of those sound possible and practical, I'll toss in one more.

Boil for some time a mixture of the alum and an excess of BaCO3, both finely powdered, together with enough water to make the mix well fluid but it need not be enough to fully dissolve the alum. The resulting precipitate of Al(OH)3, BaSO4, and excess BaCO3, should be fairly easy to filter to obtain a solution of Cs2CO3, solubility 260.5 g/100 mL and roughly 10% by weight in alcohol. The heavy barium compounds should settle rapidly, trapping the Al(OH)3. The highly soluble carbonate would seem to be a good starting point for other compounds.

blogfast25 - 21-8-2009 at 07:15

Quote: Originally posted by garage chemist

I have a german procedure for the preparation of alumina hydrate from potassium alum that uses sodium carbonate to precipitate Al(OH)3 from an alum solution. The mixed solution is kept hot for a few hours, presumably to make the Al(OH)3 more filterable.
After thorough washing, the precipitate is dissolved in HCl and re-precipitated by pouring into NH3.

Does it mention by any chance the precise pH range in which Al(OH)3 precipitates, without redissolving as Al(3+) or Al(OH)4 (-) ?

woelen - 21-8-2009 at 10:30

I tried the method, mentioned in the patent posted by Phosphor-ing. In my particular situation this by far is the easiest thing to do.

Cesium alum dissolves remarkably easily in a moderately concentrated solution of NaOH. I made such a solution and added a spatula full of cesium alum. In just a few seconds all of it dissolves. Such a solution is strongly alkaline and contains Cs(+) ions, Na(+) ions, OH(-) ions and AlO2(-) ions. The high pH allows the aluminium to go in to solution as aluminate,

My particular situation is that I have quite some NaClO4 for which I hardly had any uses (it is very hygroscopic and the material is quite wet). I dissolved some of this in water and added this solution to the solution of cesium alum plus NaOH. A few seconds after these two solutions are mixed, a compact crystalline precipitate settles. This compact precipitate is almost pure CsClO4. This precipitate, I first rinse with a solution of NaOH in order to get rid of remains of aluminate and then I rinse with water to get rid of NaOH. Then I add this crystalline solid to water, heat to boiling such that all dissolves and then let it crystallize. Now I have 1 gram of CsClO4 and the yield (relative to cesium alum) is 90%

From this CsClO4 I can make other chemicals. It is almost insoluble in cold water, but it can be converted to CsCl by simple heating. So now I have a means of making other cesium salts.

The same trick also works with NaBrO3 instead of NaClO4. I tried it. I might also work with NaClO3, but I did not try that myself. I'm not sure how well CsClO3 dissolves in water.

woelen - 21-8-2009 at 10:32

Quote: Originally posted by blogfast25

Nice purchase! Where did you buy it?

Are the solubility data yours?

The solubility data is from Wikipedia (page on cesium). The material is from eBay, it was a once only offer from a UK-based seller. This is the way I get many interesting chemicals. In the same way I also obtained 160 grams of RuO2 for just GBP 16. No repeatability, you do not know in advance what you can get, but sometimes you can be very lucky.

blogfast25 - 21-8-2009 at 12:29

Quote: Originally posted by woelen

I tried the method, mentioned in the patent posted by Phosphor-ing. In my particular situation this by far is the easiest thing to do.

Cesium alum dissolves remarkably easily in a moderately concentrated solution of NaOH. I made such a solution and added a spatula full of cesium alum. In just a few seconds all of it dissolves. Such a solution is strongly alkaline and contains Cs(+) ions, Na(+) ions, OH(-) ions and AlO2(-) ions. The high pH allows the aluminium to go in to solution as aluminate,

My particular situation is that I have quite some NaClO4 for which I hardly had any uses (it is very hygroscopic and the material is quite wet). I dissolved some of this in water and added this solution to the solution of cesium alum plus NaOH. A few seconds after these two solutions are mixed, a compact crystalline precipitate settles. This compact precipitate is almost pure CsClO4. This precipitate, I first rinse with a solution of NaOH in order to get rid of remains of aluminate and then I rinse with water to get rid of NaOH. Then I add this crystalline solid to water, heat to boiling such that all dissolves and then let it crystallize. Now I have 1 gram of CsClO4 and the yield (relative to cesium alum) is 90%

Very elegant.

How about carefully neutralising the solution of cesium alum in NaOH with fairly dilute H2SO4? The alkaline solution contains Al(OH)4 (-) (depending on concentration and pH, other aluminate species may also be present) but at about pH ≈ 7, the aluminate converts to Al(OH)3.3 H2O and if concentrated enough, will precipitate (if not, seed with freshly pure, precipitated alumina).

Now you should basically have Cs2SO4 and Na2SO4 in neutral solution. Further separate by fractional crystallisation.

len1 - 22-8-2009 at 02:37

CsClO4 is actually quite soluble in cold water, more so than KClO4 - 14.7 gm/L, so you have to use quite concentrated solutions, for a 90% yield of cesium working with 1gm quantitues, the total volume of water used must not exceed about 7ml!

blogfast25 - 22-8-2009 at 05:25

Or how about alkali digesting the cesium alum with strong ammonia, then carefully neutralise with sulphuric acid, alumina drops out, filter. Now you have Cs2SO4 + (NH4)2SO4. Crystallise and heat to decompose the ammonium sulphate to ammonium bisulphate, which is very water soluble. Selective leaching should leave the less soluble Cs2SO4 behind...

woelen - 22-8-2009 at 05:29

In water of 0 C it is just 0.8 gram per 100 ml. When excess perchlorate is used, then the solubility even goes down more. I did not measure the amount of water I used, but it is more than 7 ml (probably even double the amount, I used a test tube and it was around 2/3 filled with liquid). I did the final crystallization in water which I let cool down to 0 C (in my first attempt, I accidently had it all frozen, and I had to redo the whole operation).

woelen - 5-9-2009 at 11:59

I also tried the experiment with NaBrO3 instead of NaClO4. This also works, CsBrO3 can be precipitated from the cesium alum/NaOH solution. The yield seems to be quite good. Cesium alum dissolves in a solution of sodium hydroxide very well. One can make highly concentrated solutions. When such solutions are mixed with hot concentrated solutions of NaBrO3, then a fine white crystalline precipitate is formed when the liquid cools down.

I rinsed the fine white precipitate two times, once with dilute NaOH, in order to get rid of traces of aluminium and then with distilled water. Then I put the precipitate on fine filter paper, such that almost all liquid was sucked away from the solid. Then I transferred the solid to a test tube again and added so much water that on boiling just all of the solid dissolves. A completely clear liquid is obtained, so one can safely assume it is free of aluminate ion. I recrystallized again from this liquid and then repeated the process another time.
The final white solid was mixed with some sulphur and then it was ignited. This mix burns rapidly, like well-mixed good quality BP. It burns with a purple flame, not with a blue flame as I expected

This is an indication that still quite some sodium is left in the product. I think that the white solid I obtained is mostly CsBrO3, but I simply cannot get rid of the last traces of NaBrO3. How many recrystallizations would be needed to have a really pure product, free of sodium? Each recrystallization leads to additional loss of cesium. Apparently it is very hard to work up cesium, such that it is free of other alkalimetals. I tried with sodium hydroxide, but with potassium hydroxide I expect things to be even worse (potassium being more close to cesium).

Klute - 6-9-2009 at 10:20

Slightly related, I've recently aquired some CsCO3.. what kind of uses have cesium salts, notably in organic chemistry? The cesium carbonate major use for me in mono-N-alkylation of amines.. DO you have any planned uses for theses salts, Woelen?

woelen - 6-9-2009 at 22:58

Yes, I want to use them for experiments with light emission. Cesium ions, when excited, give a beautiful blue color and that is one of the first things which I want to explore in more depth. But this has nothing to do with organics, it is purely inorganic, probably some small pyro experiments. But up to now, my experiments have not been succesful, because the sodium ions in the material spoil the delicate spectrum of the cesium ions. Unfortunately, my method of releasing the cesium ions does not work with KOH, because KClO4 is as insoluble as CsClO4 and this does not allow me to make CsClO4 in a pure state. The same is true for KBrO3/CsBrO3. The only route is through NaOH and NaClO4 (or NaBrO3 for making CsBrO3).

garage chemist - 7-9-2009 at 01:11

You could precipitate Al(OH)3 by adding ammonia to the alum solution, filter this with suction on a buchner funnel with good paper filter, boil down the filtrate to a fraction of its volume and add HClO4 to precipitate CsClO4. This way, no sodium ions are introduced.

woelen - 7-9-2009 at 04:20

I could try this, but it sounds quite cumbersome. You need to make a LOT of solution if you want to dissolve cesium alum. Solubility is just a few grams per 100 ml of water and the cesium contents then only is about 1 gram per 100 ml of water. It can be done, but for processing 100 grams or so, I probably have to work with several liters of liquid which all has to be boiled away. Probably this is the only option left for me if I want to release the cesium without sodium contamination.

It might dissolve easily in 12% ammonia with all the cesium leaching out and getting a precipitate of Al(OH)3 in the ammonia, but there also is chance that the crystals of alum are covered by an impermeable crust of Al(OH)3 and that they do not further dissolve beyond a thin surface layer. I'll first try this on a test tube scale and will let you know about this.

[Edited on 7-9-09 by woelen]

blogfast25 - 7-9-2009 at 11:43

Quote: Originally posted by woelen

It might dissolve easily in 12% ammonia with all the cesium leaching out and getting a precipitate of Al(OH)3 in the ammonia, but there also is chance that the crystals of alum are covered by an impermeable crust of Al(OH)3 and that they do not further dissolve beyond a thin surface layer. I'll first try this on a test tube scale and will let you know about this.

No, the formation of an "impermeable crust of Al(OH)3" in those circumstances is impossible IMHO. In all likelihood the cesium alum will dissolve in excess 12 % ammonia, yielding a mixture of NH4 aluminate and Cs sulphate. Adjust pH with HCl till Al(OH)3 drops out and filter. Now you have a mix of CsSO4 and NH4Cl! Crystallize and heat to drive off NH4Cl...

woelen - 7-9-2009 at 11:53

What you are saying unfortunately is not true. Al(OH)3 does not dissolve in excess ammonia. Ammonia is not a sufficiently strong base to convert Al(OH)3 to aluminate ion. In fact, this property once was (and probably still is) used as a qualitative test for aluminium ion. Al(3+) ions give a white precipitate of Al(OH)3 when a solution of sodium hydroxide is added and this precipitate redissolves again when more sodium hydroxide is added. Al(3+) ions give a white precipitate of Al(OH)3 when ammonia is added. This precipitate does not redissolve when excess ammonia is added. This property is unique for aluminium ions. Magnesium ions also give a white precipitate with ammonia and with hydroxide, but this precipitate does not redissolve in excess hydroxide, nor in excess ammonia.

Some qualitative tests, Al and Mg are mentioned among many others:
http://www.docbrown.info/page13/ChemicalTests/ChemicalTestsc...