Derek McOlund
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Can a salt become insoluble in water?
I've been working on extracting caesium from it's mineral pollucite. After refluxing sulfuric acid and pollucite for eight hours I decanted of the
clear liquid and distilled of the water from this. Left behind in the distillation flask was conc. sulfuric acid and a white precipitated salt. From
what I've been told; it is CsAl(SO4)2 caesium alum. I tried to redissolve the alum (after decanting of almost all sulfuric acid) but nothing
dissolved.
I know that caesium alum is known for poor water solubility (along with many other caesium double salts), but how could it have dissolved in the first
place - when it later didn't dissolve at any rate?
Perhaps there is some fundamental understanding of solubilities that I don't know about that explains this?
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AJKOER
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Cesium maybe acting like iron and copper!
Fe(ll) (aq) + H+ + O2 forms Fe(OH)3. The reaction is electrochemical in nature and involves oxygen.
Similarly, basic copper(ll) salts can be formed from Cu(l) (aq) + H+ + O2 (see Wikipedia on copper oxychloride, as an example).
Boiling the clear solution in air or allowing O2 contact may be the issue.
[Edited on 14-6-2017 by AJKOER]
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DraconicAcid
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AJKOER, do you really think that cesium is oxidizing in air to form an insoluble cesium(II) or cesium(III) oxide? Or are you just trolling again?
Two possibilities- some compounds may be soluble, but be *very* slow to dissolve. Or, you may have distilled off too much solution, and the white
stuff is aluminum hydroxide or some other basic aluminum salt.
Try heating it to see if it dissolves more quickly. If that doesn't help, try re-adding sulphuric acid.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Dan Vizine
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Quote: Originally posted by Derek McOlund  | but how could it have dissolved in the first place - when it later didn't dissolve at any rate?
Perhaps there is some fundamental understanding of solubilities that I don't know about that explains this? |
The phenomenon you are referring to is much more common than you might think. It is frequently the case that during a reaction, a product or byproduct
which is insoluble in the reaction mixture is formed. It may or may not precipitate out right away. Once it does precipitate out, it won't
re-dissolve.
A perfect example of this would be the reaction of benzyl amine with tartaric acid in ethanol. If you add the solid acid to a stirred solution of the
amine and alcohol, the first thing that results is a clear solution, but the acid-base reaction has already completed. Within a few minutes a
precipitate starts to form. Once it starts, it rapidly triggers the crystallization of all the other material out of solution and there is no way to
cause it to re-dissolve in the same solvent.
[Edited on 6/15/2017 by Dan Vizine]
"All Your Children Are Poor Unfortunate Victims of Lies You Believe, a Plague Upon Your Ignorance that Keeps the Youth from the Truth They
Deserve"...F. Zappa
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elementcollector1
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If I had to take a guess, the time delay between the initial 'clear' solution and the precipitate would be due to precipitated molecules agglomerating
and forming crystals that become visible to the naked eye as they enlarge. Thus, while something might 'appear' to be soluble at first, it's actually
already precipitated out - it just hasn't clumped up yet.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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Pok
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If you let pollucite react with sulfuric acid, you will get caesium alum and silicon dioxide. The later won`t dissolve, of course. But caesium alum
will dissolve in hot (!) water. So if you heat the mixture with hot water, decant the solution and let it cool, crystals of caesium alum will form.
I did this reaction many times and never had any problems.
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AJKOER
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To be honest, my recollection confused the element Cs with Ce for which my statement is accurate. Here is an extract which I apparently was thinking
about :
"Among all rare-earth or lanthanide group elements, cerium is the only metal capable of activating H2O2 by Fenton-like mechanism. Due to its
4f26s2valence configuration, cerium is the only rare earth element to exhibit both +3 and +4 oxidation states in solution. While the cerous (Ce3+)
form is a strong reducing agent and easily oxidized by O2 in alkaline condition, the ceric (Ce4+) species is a strong oxidant under acidic condition.
Thus, cerium can easily cycle between the Ce3+ and Ce4+ oxidation states under suitable redox conditions [E0(Ce4+/Ce3+) = +1.72 V]."
Source: See "Review of iron-free Fenton-like systems for activating H2O2 in advanced oxidation processes", Link: https://www.google.com/url?sa=t&source=web&rct=j&...
-------------------------
Here is a comment on Cs found on the internet (agrees with Wikipedia, https://en.m.wikipedia.org/wiki/Caesium):
"What are the uses of Cesium?
Cesium readily combines with oxygen and is used as a getter, a material that combines with and removes trace gases from vacuum tubes. Cesium is also
used in atomic clocks, in photoelectric cells and as a catalyst in the hydrogenation of certain organic compounds."
Per Atomistry (see http://caesium.atomistry.com/caesium_monoxide.html ) even the sub oxides readily react with air!
Another source (http://www.lenntech.com/periodic/elements/cs.htm ) notes that Cs is the most electropositive of all elements! So why is my statement on possible
electrochemical based reactions with air attacked?
--------------------------------------------------
My revised comment is that not only mainly Cs but any other transition metals impurity present (from say ore processing) with air and any light
exposure likely contribute to an electrochemical or possible Fenton-type reactions producing hydroxyl radicals and OH-. The latter directly
contributes to the creation of basic salts. Any hydroxyl radicals also formed can contribute to reactive oxygen species leading to a partial recycling
of a lower valence state that can feed Fenton-type reactions (more likely agents here lowering valence states are sunlight and the maintenance of
equilibrium relationships in any formed redox couples).
[Edited on 21-6-2017 by AJKOER]
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