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Author: Subject: Turning nitrates into Ca or Mg nitrate?
Refinery
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[*] posted on 21-3-2014 at 08:19
Turning nitrates into Ca or Mg nitrate?


I know that when mixing ammonium nitrate solution with calcium hydroxide and heating ammonia gas is evolved and calcium nitrate is formed.

But how we can make calcium or magnesium nitrate from potassium or sodium nitrate? I know sulfuric acid yields nitric acid with any nitrate, but this method is not an option in this case. Could I use some metathesis or double displacement reactions?
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[*] posted on 21-3-2014 at 12:38


Yes it is possible, but not very cost effective. What you would need to do is first turn the potassium nitrate or sodium nitrate into ammonium nitrate by first determining the moles of nitrate present, make it into a concentrated solution, and make a concentrated solution containing the same number of moles of sodium bisulfate. Next add ammonia until the solution is basic, then boil until a few crystals can be seen. Chill the solution to 0oC and filter. Evaporate the filtrate to get ammonium nitrate. You can now mix this with a hydroxide and heat to get the desired nitrate and ammonia gas if done in a concentrated solution. If you don't mind the nitrate being impure or a low yield, you can check out the solubility tables. Through a metathesis like this one, only a soluble group II salt can be used, and you must ensure that the solubility of one of the products is lower than the solubility of both reactants in the solvent you are using. Then a metathesis will work, but it is not preferable in a situation like this where all of the products and reactants will be soluble in water.
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[*] posted on 21-3-2014 at 13:31


First could you tell me what you want to do with your product? This can help me figure out what the best method(s) would be.



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[*] posted on 21-3-2014 at 14:20


Quote: Originally posted by Refinery  
But how we can make calcium or magnesium nitrate from potassium or sodium nitrate?


The simple truth is that almost any nitrate is made from nitric acid: N2 + H2 to NH3 (Haber), then NH3 to HNO3 (Ostwald). Then HNO3 to any nitrate.

Complicated schemes like gdflp's are expensive and often lead to very impure products. I've seen countless of them proposed here but almost none ever followed up because they are just not practical.

And what would you want calcium nitrate for?

[Edited on 21-3-2014 by blogfast25]




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[*] posted on 21-3-2014 at 14:50


blogfast25's answer is by far the most accurate. If you want a pure product, nitric acid is the way to go. If you don't own nitric acid you can make it but I wouldn't recommend it because you seem to be at an entry level of skill. If you do feel that you can produce nitric acid and have the right stuff it should be pretty simple. Our very own UnintentionalChaos has great video on producing ~90% nitric acid. If you follow his instructions and have your nitric acid it's a simple matter of adding the acid to magnesium or calcium hydroxide. To my knowledge it will be very difficult to crash the nitrate salt out of solution so you will probably be forced to boil it down to a very low volume where you can see crystals forming in large amounts and then letting it evaporate from that point on.



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[*] posted on 22-3-2014 at 05:21


And you don't need 90 % for most nitrates: even 35 % will dissolve most things that 90 % will, only slower.

The one time I boiled down a solution of CaCO3 in nitric acid I got a very hygroscopic product, fairly useless really...




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[*] posted on 22-3-2014 at 05:43


You may wish to try this more expense route that I have performed. You start with KNO3 which I obtained cheaply as a stump remover sold in Home Depot (99% pure for around $8.00). Then mix in some Epsom Salt dissolved in hot distilled water:

KNO3 + 2 MgSO4 = K2SO4 + Mg(NO3)2

On freezing the mixture, long cactus like needles of K2SO4 separate out. Filter, done, Magnesium nitrate. Keep the Potassium sulfate, a good source of K for making KClO3.

If one adds dilute aqueous ammonia to your aqueous Magnesium nitrate, a white solid (Mg(OH)2) precipitates leaving largely aqueous NH4NO3. To this clear filtered solution if you add more dilute ammonia, yet more Mg(OH)2 appears. Filter again and repeat as needed for your desired purity of NH4NO3.

[Edited on 22-3-2014 by AJKOER]
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[*] posted on 22-3-2014 at 09:31


Quote: Originally posted by AJKOER  
On freezing the mixture, long cactus like needles of K2SO4 separate out. Filter, done, Magnesium nitrate. Keep the Potassium sulfate, a good source of K for making KClO3.



Even at 0 C, the solubility of K2SO4 is still about 7.4 g / 100 g of water. Your magnesium nitrate is contaminated.

Since as Mg(NO3)2 is so ridiculously water soluble, you could boil down the solution until crystals of K2SO4 start forming, then chill, thereby reducing the K2SO4 content further. Repeat till you're blue in the face.

Not worth doing, IMHO...




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[*] posted on 22-3-2014 at 09:42


Well, a few recrystalizations should make it pure, I still might give it a try...
Potassium nitrate is so cheap, it's worth it for me.




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[*] posted on 22-3-2014 at 12:09


OP you just don't have a lot of viable options here. Could you substitute the calcium or magnesium nitrates with other nitrate salts like potassium?



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[*] posted on 22-3-2014 at 12:28


You make it from dollars not alkali nitrates! I bought a 40lb sack of calcium nitrate from the farm supply. Dissolve, filter from the sodium/calcium stearate in it, and you're done.

Much more practical. Currently I'm attempting to grow a calcium phosphate tube by diffusion through cardboard, Ca(NO3)2 outside and Na3PO4 inside the tube, fixed into a plastic cup with some wax. What fun.
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[*] posted on 22-3-2014 at 16:43


I want to produce NO2 gas by thermal decomposition of Ca or Mg nitrate, that would be the only purity requirement. If it has sulfates, carbonates or oxides in it, it won't matter, but if it contains potassium, sodium or ammonium nitrate, it will be a huge safety risk because the mixture is heated over 650C. Mg nitrate will decompose at much lower rate, but Mg carbonate, oxide and hydroxide are a lot more expensive to me, like 10 to 30 times.

I find only mixed fertilizers in my local stores, but I checked the contents and they should up mostly ammonium nitrate, urea and some chlorides and calcium carbonate. I found that mostly all fertilizers contain just ammonium nitrate, very few have others like potassium or calcium, I dont know why. This way I think I can just add Ca(OH)2 to it and boil it up to get Ca nitrate, Ca carbonate(from urea, I think) and other impurities, then filter it and then concentrate with some extra filtering until it's mostly calcium.
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[*] posted on 23-3-2014 at 05:06


Quote: Originally posted by blogfast25  
Quote: Originally posted by AJKOER  
On freezing the mixture, long cactus like needles of K2SO4 separate out. Filter, done, Magnesium nitrate. Keep the Potassium sulfate, a good source of K for making KClO3.



Even at 0 C, the solubility of K2SO4 is still about 7.4 g / 100 g of water. Your magnesium nitrate is contaminated.

Since as Mg(NO3)2 is so ridiculously water soluble, you could boil down the solution until crystals of K2SO4 start forming, then chill, thereby reducing the K2SO4 content further. Repeat till you're blue in the face.

Not worth doing, IMHO...


A good point if the solution was made dilute. However, the process starts with two dry salts, KNO3 and dehydrated Epsom's Salt. So only add sufficient boiling distilled water to dissolve the KNO3 and MgSO4. This should largely reduce any K2SO4 presence, but as Blogfast stresses, be mindful of contamination issues.
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[*] posted on 23-3-2014 at 05:32


If you want nitrogen dioxide, a much cheaper way IMHO is instead react copper, a nitrate salt and hydrochloric acid. It doesn't need high temperatures and NurdRage has a great video showing the process here http://www.youtube.com/watch?v=2yE7v4wkuZU
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[*] posted on 23-3-2014 at 06:06


Quote: Originally posted by gdflp  
If you want nitrogen dioxide, a much cheaper way IMHO is instead react copper, a nitrate salt and hydrochloric acid. It doesn't need high temperatures and NurdRage has a great video showing the process here http://www.youtube.com/watch?v=2yE7v4wkuZU


No, although chemically interesting, this is another silly way to destroy a valuable nitrate to... obtain a nitrate.

Get distilling, for Science sake!

I generally like Nurdrage but there's nothing 'complete' about this guide...



[Edited on 23-3-2014 by blogfast25]




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[*] posted on 23-3-2014 at 06:14


Quote: Originally posted by AJKOER  

A good point if the solution was made dilute. However, the process starts with two dry salts, KNO3 and dehydrated Epsom's Salt. So only add sufficient boiling distilled water to dissolve the KNO3 and MgSO4. This should largely reduce any K2SO4 presence, but as Blogfast stresses, be mindful of contamination issues.


It makes not a blinding bit of difference whether you start from a dilute, a concentrated or what not solution: after chilling if you get K2SO4 crystals then the supernatant is by definition still saturated with K2SO4. Of course if the volume of supernatant is small, then total amount of K2SO4 therein contained will be small. But it's there. Always.

[Edited on 23-3-2014 by blogfast25]




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[*] posted on 23-3-2014 at 06:38


No my point was that he wanted calcium nitrate to make nitrogen dioxide, and using hydrochloric acid, copper, and a nitrate is a much safer way to make nitrogen dioxide than heating a nitrate to decomposition.
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[*] posted on 23-3-2014 at 07:13


Quote: Originally posted by blogfast25  
...

It makes not a blinding bit of difference whether you start from a dilute, a concentrated or what not solution: after chilling if you get K2SO4 crystals then the supernatant is by definition still saturated with K2SO4. Of course if the volume of supernatant is small, then total amount of K2SO4 therein contained will be small. But it's there. Always.


In cases where purity is a concern, one may be able to do better than suggested by Blogfast's constant percentage contamination argument on the smallest practical volume.

The more laborious process would start with say an stoichimetric excess of KNO3. Then, a common ion effect (from the Potassium in the excess KNO3) would reduce the relative solubility of the unwanted K2SO4. Cool and separate out the K2SO4. Note, one could repeat the process many times by restocking the solution with MgSO4 maintaining an excess of KNO3. Then add the required MgSO4 to remove the excess KNO3 and boil down the solution and collect the K2SO4.

I would argue that the final solution may have a lower contamination level following this process. Note during the process, the supernatant is still concentrated with Potassium cations, but all of them are not associated with sulfate anions.

[Edited on 23-3-2014 by AJKOER]
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[*] posted on 25-3-2014 at 07:54


Hmm.. It seems that there is another problem. The ammonium nitrate is mixed with potassium sulfate, which both are well water soluble. As far as I thought it, if I mix calcium hydroxide into there, the calcium nitrate formed will react with the potassium sulfate to form potassium nitrate and calcium sulfate, which precipitates out. The potassium sulfate could be separated with methanol, though. If there is substantial amount of KNO3 in Ca(NO3)2 when heated to 650C, it may cause major risk.

The solubility of AN in boiling water is about 10kg per liter, whereas potassium sulfate will do only 240g per liter. Probably it will salt out far before this, because potassium sulfate is likely not soluble in AN.. :)

[Edited on 25-3-2014 by Refinery]
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[*] posted on 25-3-2014 at 13:38


Personal observation, if one mixes NaCl and MgSO4, the reaction:

CaCl2 + MgSO4 = CaSO4 + MgCl2

does not move largely to the right with a significant precipitation of CaSO4. Apparently, Calcium sulfate is soluble under acidic conditions.
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[*] posted on 5-8-2015 at 18:49


In Handbook of Preparative Inorganic Chemistry, Baur suggest a very simular metasynthesis for the ammonium chlorate intermediate in his barium chlorate preparation. The technique involves mixing the potassium chlorate solution with ammonium sulfate and boiling down until a slurry forms. Next an alcohol solution of 80% strength is added to help the potassium sulfate crash out more completely. Ammonium chlorate is only moderately soluble in alcohol. This should work better for magnesium nitrate. With lithium salts, you could use acetone.

Ex(LiCl + KNO3 >>>acetone soln>>KCl +LiNO3




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