I decided to put this in Miscellaneous because it’s Phys. Chem and also unlikely to interest those whose life revolves around p and sp2 orbitals,
hairy neucleophilic radicals, etc…
The following are examples of unusual or difficult separations.
(1) Rare Earth Sulphates
I recently did two experiments to get rare earth salts from common materials (lighter flints (Ce) and magnets (Nd)) (qv thread via search engine). I
will shortly add Sm to these and sacrifice a SmCo5 magnet (and I won’t throw the Co away, you bet!)
In these cases the sulphates are much less soluble in hot water than cold, ~1g/100g aq. @ 100C. If you have spent as many efforts as I have
crystallizing compounds by cooling the solution, it seems very odd to have to keep the solution as hot as possible for maximum yield. And washing with
boiling water seems wrong, too.
NOTE: if anyone has the solubility/temperature for samarium sulphate, I’d appreciate any data. I have data for most RE but not Sm sulphate.
Another compound that exhibits similar trends is LiCO3, but not as radically. Many sulphates, in fact, tend to have a peak water solubility somewhere
below 100C.
(2)NaCl.2H2O, an unstable hydrate.
In recent efforts to produce permanganates (see thread, if interested), I was faced with the problem of extracting the permanganate ion from a mixture
of sodium chloride (lots of it), sodium chlorate, and carbonate (much less of those two), and sodium permanganate (lowest, far less than I’d
hoped!). The solubilities are in order, highest to lowest:
NaMnO4>NaClO3>NaCl~NaCO3.
The solubility of the permanganate is very high, that of chlorate high, chloride essentially flat with temperature ~ 37g/100g aq., and carbonate is
far less at 0C ( ~7g/100g aq.). In the presence of excess Na+ ions, it’s solubility is probably a lot lower than that (common ion effect). Hence the
carbonate can be separated to some extent by cooling, after first saturating the mixed solution by evaporation until a good deal of the NaCl
separates. Then most of the carbonate and a little more chloride will be precipitated on cooling.
The normal procedure would be to further evaporate and precipitate NaCl progressively. When the liquid, now containing permanaganate, chlorate and
chloride and a bit of carbonate gets concentrated, it becomes increasingly difficult to filter off the solid (through glass frit or wool, because of
permanganate). I decided to cool it. Nothing precipitated down to -10C. I assumed the NaCl was supersaturated and left it cooling.
A mass of crystals separated somewhere between -10C and -20C. They were not ice, which would have floated; the chlorate and permanganate was still
assumed to be highly soluble and it seemed unlikely to be them. Nor carbonate. So these crystals contained NaCl, but instead of being neatly cubic,
they were thin flat rhomboidal plates.
A bit of research revealed that they must be NaCl.2H2O, a hydrate that is only stable between -23C (the eutectic of NaCl & H2O) and 0C, where they
decompose into cubic NaCl, and solution.
Was this a neat way of removing NaCl? I then did an experiment using pure NaCl and water, producing a saturated solution and repeating the same
cooling cycle. I used 100g (cc) aq as solvent. Collecting the NaCl.2H20 crystals carefully (difficult) I let them warm up, when a solution of NaCl
and the expected cubic crystals appeared. I evaporated the water off after warming, and found they contained 9.90g of NaCl. Since originally there
were 36.1 g in a saturated solution with 100 g aq., this process removes about 27% 0f the NaCl from a saturated solution. Not spectacular but a help.
(3)Using organic solvents for inorganic salts.
Following up on the separation attempted in (2) above, one finally obtains a solution containing mainly chlorate and chloride with the
permanganate a lesser component. Sodium permanganate is by far the most soluble (approx >200g/100g aq. versus about 100/100 for chlorate, with
chloride at about 36/100 in isolated solution). Thus ultimately chlorate will be in excess, and a portion of the chlorate can be precipitated. But at
this point the solution is very concentrated and also of high viscosity and very difficult to filter or decant effectively.
A quantity of KCl is then added to precipitate the much less soluble KMnO4, sol. 2.8g/100g at 0C. Unfortunately the sol. of KClO3 is also low,
3.3g/100g.
A mixed precipitate of permanganate and chlorate is obtained. The two salts are not isomorphous (perchlorate is) but an intimate mixture is obtained.
In order to precipitate all the desired permanganate, an excess of KCl is needed and all the chlorate must be precipitated too. The separation of
salts so close in solubility would be a Herculean task. Instead, the solution is evaporated to dryness.
There are a few organic solvents capable of dissolving some inorganic salts. Generally they must possess a high dielectric constant (>~15) and
hence be polar. Among these are alcohols, ketones, nitriles and amides. Alcohols react with KMnO4 and are unsuitable. The commonest ketone, acetone,
does dissolve KMnO4 (~10g/100g acetone) and does not dissolve KClO3 significantly, nor NaCl or KCl.
However, if significant amounts of the acetone are in the enol form the same reaction is possible as with alcohols. To avoid this I keep the
temperature low and make sure the solvent is dry. CaSO4 is the best agent but it doesn't take up much water.
It works but is not a very practical method. I evaporated the acetone in a stream of air from a fan (outside! Very inflammable, low flash point). Not
a good idea; acetone is very hygroscopic and easily picks up water during the evaporation. Especially in high humidity. A lot of acetone is needed for
a little KMnO4. But extraction can be achieved this way.
A decent setup would require exclusion of humid air and recovery of the solvent. Probably best done under partial vacuum.
I was flabbergast (great word) to read about the NaCl.2H20, I have never come across a reference to that before. Such unusual behaviour from such a
simple household chemical. It reveals there is still much to be discovered, even by amateurs.
Re. KMnO4. I guess I should have tried acetone extraction and evaporation when I was making it by electrolysis. It would have resulted in a higher
yield, but as you point out a decent recovery system is really required!
Regards, XenoidDerAlte - 27-10-2007 at 12:28
@ Xenoid
NaCl.2H2O might be completely unknown to today's amateurs but the old chemists in the 1800's knew all about it!
Re KMnO4, I don't think you would have had to use acetone to extract more KMnO4. You do not have the same problem, which is caused by the use of
hypochlorite as an oxidant and the close solubility of KMnO4 and KCl03. In your process the likely products are remaining MnO2, NaOH, and nitrite
probably reconverted to nitrate by anodic oxidation, as well as the desired KMnO4 and maybe some K2MnO4. Plus some NaCl and KCl from the added
potassium salt. (Actually, of course, just a mixture of ions!)
Nitrite and nitrate are quite soluble. MnO2 can be filtered off. Evaporation and cooling to 0C should do the trick of depositing most of the
permanganate.
I used commercial acetone solvent. It seemed pretty pure except for a bit of water. But, as I pointed out, it's hygroscopic. And it does not seem to
be too stable with KMnO4 once there's water around, probably due to enol formation. To do it well is a bit of a chore. And even commercial acetone is
nothing like as cheap as water (even if you are stupid enough to buy it in bottles!).
