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Author: Subject: Ammonium nitrate purification
ScienceSquirrel
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[*] posted on 1-3-2013 at 17:10


Quote: Originally posted by DraconicAcid  
Quote: Originally posted by ScienceSquirrel  
Calcium hydroxide has a solubility of 0.173g/100ml at 20C and a pKa of 12.4.
It is a stronge but poorly soluble base.
Barium carbonate has a solubility of 0.0024g/100ml and a pKa of maybe 8 or 9.

I think you mean that the pH of the saturated solutions are 12.4 and maybe 8 or 9. As they are not acidic in aqueous solutions, they don't have pKa values. (Calcium hydroxide could act as an acid in a non-aqueous solvent and a blisteringly strong base, such as hydride ion, but that's not relevant to this discussion.)


Calcium hydroxide displaces ammonia from it's salts in aqueous solution or suspension.
Calcium or barium carbonate does not.
These are relevant facts to the discussion.


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AJKOER
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[*] posted on 1-3-2013 at 17:30


Here is a reference (source: "A Laboratory Manual of Qualitative Chemical Analysis, for Students of ...", by Andrew Richard Bliss, page 122, link: http://books.google.com/books?pg=PA122&dq=baco3+soluble+... ) that is nearly precisely the same as the comments by Science Lab on BaCO3. To quote:

"TESTS [for Barium]
324. Ammonium Carbonate. (NH4)2CO3, and soluble carbonates produce white precipitates of Barium Carbonate, BaCO3. If the solution is acid, complete precipitation takes place only after heating. The precipitate is easily soluble in diluted Hydrochloric Acid, HCl, in Nitric Acid, HNO3, and Acetic Acid, HC2H3C2. Slightly soluble in excess of Ammonium Chloride, NH4Cl."

This source ("Solubilities of Inorganic and Organic Compounds: A Compilation of ..", by Atherton Seidell, page 108, link: http://books.google.com/books?pg=PA108&id=d1JMAAAAMAAJ#v... ) states: "Barium carbonate boiled with aqueous NH4Cl is slowly but completely decomposed. The time required varies inversely as the concentration of the NH4Cl solution." However, if dissolving the BaCO3 appears to be a problem, add some HNO3 to the NH4NO3 solution mix (or, add a little Oxalic acid to form some HNO3 in situ before adding the BaCO3).

Now as to why, here is another educational reference (Question 59) under the heading of "Dissolution of Precipitates and Complex Ion Formation" that asks the student to explain why any of Mg(OH)2, Mn(OH)2 or Ni(OH)2, can be dissolved upon the addition of NH4NO3 or NH4Cl solution. The answer appears to relate to complex ion formation. Reference: page 792 at http://books.google.com/books?id=6Zwu9-qT0qQC&pg=PA792&a...


[Edited on 2-3-2013 by AJKOER]
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[*] posted on 7-3-2013 at 16:18


Sorry, AJKOER. I have no barium compounds available. I also got some problems to acquire my glassware. I am going to wait for a while, until I finally get all of them. I have been very busy lately. I will keep doing my experiments next month.



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[*] posted on 25-10-2020 at 08:41


Could dual solvent(EtOH + H2O) be used to reduce solubility of AN to allow for easier recrystallization?

But then, the residue will have chlorides, and it will generate hydrochloric acid when mixed with H2SO4. What if we turn things around, and instead bubble HCl directly into the concentrated solution to generate HNO3 and make chlorides out of everything else? HCl could be losslessly generated from common salt with H2SO4 as gas.

[Edited on 25-10-2020 by Fyndium]

[Edited on 25-10-2020 by Fyndium]
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[*] posted on 1-11-2020 at 15:24


An example of yield for KNO3 recrystallization.

A batch of 2000g of source that contains 50% KNO3 according to MSDS, was processed. It was first dissolved in about 50% weight of water, and heated until no more dissolved. During process it appeared that it contained something that dissolves at lower temps, but precipitates upon heating. The solution was filtered hot, and clear, but deeply dark colored liquid obtained. It was left for recrystallize until ntp, and then it was transferred to cool and cooled down to around 2-4C, decanted and suctioned dry. Crystal mass, seemingly containing co-precipitates with signs of KNO3 crystals was obtained.

This mass was dissolved in about 50% of mass of water and heated until clear solution was obtained and it was filtered again. It was left for crystallize, and large, multiple cm long crystalline shards formed, and after ntp it was cooled down to near zero, decanted, suctioned dry, washed with minimum amount of ice water and rinsed with little ethanol and dried over steel mesh with fan. Yield of dry crystalline mass was 640g.

The yield appears low, at 64%, but considering that minimum of 14g/L of KNO3 will remain dissolved in the mother liquor in two steps, and considering the washing and rinsing losses, the realistically achievable yield can be calculated even over 90%. The large mass of impurities requires larger than minimum amount of solvent for recrystallization, because lower volumes causes co-precipitation.

Considering that this source is otc beyond description, cheapish* and results in useful amounts of highly pure crystalline reagent, I see it as a viable method. KNO3 is probably the easiest nitrate to refine from otc sources, and at least conceptually, it can be metathesized from any other nitrate by adding potassium chloride into the solution, which will cause KNO3 to precipitate upon cooling, it being the least soluble molecule. Carbonates in case of calcium, magnesium etc could be precipitated out the other way, but solubilities of Na and K carbonates are too close to KNO3 so they are not preferred. KCl itself is easily recrystallized pure from otc salt-free salt.

*expensive compared to farming stores which sell out 25-40kg bags of practically pure KNO3, but they generally are aimed for professional market and it is much possible that an individual purchasing a bag with cash could trigger some red flags somewhere because high nitrogen fertilizers are officially listed as monitored substances in multiple countries due to the grave potential for misuse - and very importantly, hobbyists generally don't needs tens of kg's of reagents for their experiments.
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