Sciencemadness Discussion Board - Ammonium nitrate purification

Ammonium nitrate purification

dulio - 24-2-2013 at 15:45

Hi, folks. Beforehand, I would like to point what I have read so far:

http://www.sciencemadness.org/talk/viewthread.php?tid=1112

http://www.sciencemadness.org/talk/viewthread.php?tid=21362

http://www.sciencemadness.org/talk/viewthread.php?tid=3657

http://www.sciencemadness.org/talk/viewthread.php?tid=1058

My issue is simple. I have got some ammonium nitrate and I want to purify it. The manufacturer claims that the fertiliser contains 32% of nitrogen but several impurities are present. The label also says it contains 1% of potassium oxide. No more information is provided. Typical of poor brands like the one I have found.

The fertiliser consists of a heterogeneous mixture of four components. The major one is ammonium nitrate itself, small white granules. I gess their diameter is ~5 mm. Potassium oxide is pretty similar, but it is yellow. Pinkish rocks, bigger than the white granules, are present together with dark granules. A sort of phosphates, I suppose. If I am not wrong, there are more pink nuggets than dark ones. The sample is not here right now.

I do not know what the pink rocks are, but maybe it is calcium carbonate. Oh, I have so many questions... Anyway, let us start from the basics. How can I purify the nitrate in order to further convert it into sodium nitrate? I also want to keep some pure stuff in my shelf.

Any help is welcome. Thanks in advance, people.

elementcollector1 - 24-2-2013 at 16:09

I would suggest screening out the larger particles, allowing the ammonium nitrate and "potassium oxide" (which is probably really potash, or even garden sulfur) to pass through. Then, if the stuff really is potash and not sulfur, I would suggest recrystallization from boiling water.

cyanureeves - 24-2-2013 at 16:11

i too have about 3lbs. of what i think is urea in my ammonium nitrate that makes my sodium nitrate bubble like mad.i dont think i will save this mixed batch because i've lost alot of sulfuric acid when making nitric acid.dollar general ice compress really let me down and i will never amass any of their nitrate.i never stopped to think why i didnt see sodium nitrate crystals forming as before even though urea also makes nice shards all i get is doughy paste.

jock88 - 24-2-2013 at 16:32

Funny I had bags of Calcium Ammonium Nitrate fertilizer and lots of pink 'balls' (prills) were in with the usual balls of Ca ammonium nitrate. Never seen them before.
Methonol is used to purify Ammonium Nitrate by recrystillization.

BromicAcid - 24-2-2013 at 16:41

Basic information needed. Where did you get the information regarding the components of the mixture? Did you manage to find the MSDS? Or was this information from the NPK number where your bag would be saying 32-0-1? If it is from the NPK number and it is 32-0-1 then you have no phosphate and can ignore that. Also your potassium is NOT in the form of K2O, the NPK number is an older system and simply tells you what it is in terms of it's component % if it were dead burnt. It could be in the form of the nitrate as well or the phosphate (if there was P in the NPK number) or silicate or any of a number of other ions.

Pure NH4NO3 has a NPK number of 33-0-0 so you would be pretty close to the real stuff. What is your final use for the material and what final purity are you hoping to obtain? Your material might be >95% pure which would be in the reagent grade area. Also, it is folly to assume that the different colored and textured bits in your fertilizer are different materials. Sometimes they 'cement' highly soluble materials like ammonium nitrate with clay or the like to make them 'slow release' or in the case of ammonium nitrate to make them less likely to function as an oxidizer. There might also be retardants present to the same purpose such as the more widely known flammability retardants present in sulfur sold for landscaping purposes.

If you want pure material you will likely need to recryst the material anyway. That is where you should focus your efforts or at least that is where I would focus if I were you and I needed fairly pure material. Wiki states that about a kilo of the material will dissolve in boiling water but only about 10% of that at 0C, that's decent recovery to get rid of insoluble garbage and get some nice crystalline material of better purity. Note though unless you start dilute and boil it down you will need to do a hot filtration which gets nice and tricky when dealing with concentrated solutions.

dulio - 24-2-2013 at 16:51

Well, elementcollector1, it wound not work. The mixture is very wet and the components have similar size. I have bought two pairs of tweezers and separated half of the package content (which contains 1 kg). It took me a lot of time but I wish to compare the results using both the raw material and the separated one. My brother is not at home now. We he arrives, I will post the pictures I have taken. It will make things clearer.

I believe it is potassium oxide rather than sulphur because it is specified in the label. Anyway, I am not worried about it because, since it is present only in a very small amount, I can easily remove the yellows rocks just picking them out with the tweezers.

The dark rocks are pretty much similar to these:

What is that? Regular phosphate, I mean PO₄^3-.

Does anyone have information on the solubilities of ammonium nitrate in ethanol? What about the phosphates? I have some ideas in mind but, first of all, I would like to identify the components to try something really effective. Mistakes must be avoided at all cost.

dulio - 24-2-2013 at 17:05

Well, HBr. The 32% is on the box. There is no NPK number. As I told, it is a poor brand. I could not find any other ammonium nitrate source. I will keep looking for.

The 1% of potassium oxide is also in the box. Such small amount fits the physical evidence of only a very few small yellow rocks within the mixture.

The mixture seems too heterogeneous to be that pure. I live in Brazil, man. We do not trust labels here. At least I do not. A few years ago a chemical engineer was arrested for putting sodium hydroxide and hydrogen peroxide in milk.

How should I start? I wonder if I have to recrystalize twice. First with water and after with methanol.

PS.: By the time I posted, BromicAcid have not answered yet.

dulio - 24-2-2013 at 17:16

I forgot to tell why I need that nitrate. I want to have some pure chemical in my shelf. I also want to turn it into sodium nitrate reacting it with sodium carbonate. I will also make another experiment with sodium bicarbonate, just to decide which one is better. I think the carbonate will be preferable. I should also mention that I wish to collect the released ammonia in deionised water.

What about you, jock88? Do you think you have got the same impurity? The pictures will make it clearer, but I would like to anticipate that most of the "prills" are pinkish whilst some of them are white. All the components are hygroscopic.

BromicAcid - 24-2-2013 at 20:25

@ dulio

Pure ammonium nitrate has a NPK number of 33 - 0 - 0, i.e, 33% nitrogen. It is because your number is so close to this that I expect you have fairly pure material and the numbers on the box are the NPK numbers, however in this case there is no phosphate so all they are reporting is the nitrogen and the potassium 32 - 0 - 1

However... if all you have is % nitrogen and it does not specify this as ammonium nitrate then you might have a mess. Urea is 46-0-0 and is used quite a bit as a fertilizer too. So this could be a urea based fertilizer if it did not specificy ammonium nitrate and if that is the case there's plenty of room for other crap in there, i.e., iron sulfate, sulfur, and the other things they add to fertilizers for plants that wouldn't show up as NPK components.

I would still hold out hope though that you have a mostly pure product. Your approach seems reasonable to separate out the different components by look and then try to work them up. Do some small scale work to see if when you dissolve it if you end up with insolubles and if they sink or float. Best case scenario would be being able to dissolve in hot water and decant away from solids avoiding the hot filtration. The solubility in water drops off quickly with temperature so the possibility of blinding your filter paper and forcing yourself to deal with near boiling water is a real possibility.

dulio - 24-2-2013 at 21:24

Actually, I heard it is ammonium nitrate. The manufacturer only says that it contains 32% nitrogen and 1% potassium oxide. Maybe I will carry on with my investigations this Thursday. I should post pictures as soon as I can.

blogfast25 - 25-2-2013 at 13:06

Quote: Originally posted by BromicAcid

If you want pure material you will likely need to recryst the material anyway. That is where you should focus your efforts or at least that is where I would focus if I were you and I needed fairly pure material.

Hmmm… Wiki actually states the following solubility data for ammonium nitrate (water):

118 g/100 ml (0 °C)
150 g/100 ml (20 °C)
297 g/100 ml (40 °C)
410 g/100 ml (60 °C)
576 g/100 ml (80 °C)
1024 g/100 ml (100 °C

If correct, then that stuff is so water-soluble that, despite the strong temperature-solubility gradient, it would be very hard to recrystallise it by cooling a hot, concentrated solution. The entry further states:

”For industrial production, this is done using anhydrous ammonia gas and concentrated nitric acid. This reaction is violent and very exothermic. After the solution is formed, typically at about 83% concentration, the excess water is evaporated to an ammonium nitrate (AN) content of 95% to 99.9% concentration (AN melt), depending on grade. The AN melt is then made into "prills" or small beads in a spray tower, or into granules by spraying and tumbling in a rotating drum. The prills or granules may be further dried, cooled, and then coated to prevent caking. These prills or granules are the typical AN products in commerce.”

Even 410 g/100 ml (at 60 C) is more a solution of water in molten ammonium nitrate! A 50 % solution at say 100 C would still not yield any crystalline material when cooled to 0 C…

[Edited on 25-2-2013 by blogfast25]

elementcollector1 - 25-2-2013 at 13:12

Still, chilling a saturated solution of AN at 100 C to 0 C causes about 9/10 of the AN to precipitate out...

blogfast25 - 25-2-2013 at 13:40

Quote: Originally posted by elementcollector1

Still, chilling a saturated solution of AN at 100 C to 0 C causes about 9/10 of the AN to precipitate out...

Ahem. A saturated solution of AN at 100 C is 1024 g of AN per 100 ml of water (you still call that a solution??) I call that 'molten AN with 10 % water contamination'...

Trust me, thermally recrystallising AN is NOT an option here. For a proper recrystallisation you need some water so your impurities can gather there. In the case of AN everyting will simply solidify, crap included.

[Edited on 25-2-2013 by blogfast25]

[Edited on 25-2-2013 by blogfast25]

AJKOER - 25-2-2013 at 20:44

Here is a suggestion that others may speak well of (unlikely I suspect). Nevertheless, here it is:

1. Dissolve your AN mixture in dilute Oxalic acid.

2. The solution remaining after many of the oxalates have precipitated out should contain mainly dilute HNO3.

3. React with Cu to form NO to combine with O2 to form NO2.

4. Dissolve the NO2 in water and treat with more O2 to form pure HNO3.

5. React with NaOH to prepare NaNO3.

weiming1998 - 26-2-2013 at 00:59

Quote: Originally posted by AJKOER

Here is a suggestion that others may speak well of (unlikely I suspect). Nevertheless, here it is:

1. Dissolve your AN mixture in dilute Oxalic acid.

2. The solution remaining after many of the oxalates have precipitated out should contain mainly dilute HNO3.

3. React with Cu to form NO to combine with O2 to form NO2.

4. Dissolve the NO2 in water and treat with more O2 to form pure HNO3.

5. React with NaOH to prepare NaNO3.

If the nitrate ions are going to be converted to nitrogen oxides anyway, there is no point in precipitating any impurities. Simply dissolve the impure fertiliser in excess concentrated hydrochloric acid, add copper and bubble the resulting gases through water will make dilute HNO3 that can be used to make NaNO3.

[Edited on 26-2-2013 by weiming1998]

AJKOER - 26-2-2013 at 08:44

Weiming1998:

Yes, you are right, but one could stop the process before adding Copper if impure HNO3 (mixed with some phosphoric acid, etc.) was acceptable.

Also, the H2C2O4 may more favorably react with (or decompose) organic compounds upon warming (for example Urea).

woelen - 26-2-2013 at 10:39

The normal way of obtaining ammonium nitrate of decent purity is to dissolve all of the impure ammonium nitrate based fertilizer in water. Take 5 liters of water and add 5 kg of the fertilizer. Stir every few minutes, until all granules are broken apart and you have a fairly homogeneous turbid liquid. Allow the solution to settle overnight. A layer of insoluble crap collects at the bottom of the bucket. Using a PVC tube, take away most of the clear liquid (e.g. 8 liter can be recovered easily). The remaining 2 liter can perfectly be used as fertilizer for your garden (it contains quite some calcium carbonate, remains of ammonium nitrate and other fertilizer-worthy insoluble stuff). Just spreay it around on a rainy day to make your plants happy.

The 8 liters of liquid must be allowed to dry. Pour some of it in a dish and put that on a hot radiator and allow it to stand there for a day or so. Finally you will get crystals, but these crystals will remain somewhat humid. It is hard to obtain a really nice and dry product. You could try drying in an oven, but do not overheat the product (it easily decomposes). Heating to 80 C or so can be done safely though. In this way you can dry several batches.

You could also add potassium carbonate to the clear liquid to get potassium nitrate, which is much easier to purify. Do this reaction outside! A LOT of ammonia fumes will be produced!

[Edited on 26-2-13 by woelen]

blogfast25 - 26-2-2013 at 10:51

If you really want to convert the crude ammonium nitrate to sodium nitrate, react moistened crude AN with a stoichiometric amount of solid NaOH, heating slightly will help:

NaOH + NH4NO3 == > NaNO3 + NH3(g) + H2O (Classic Displacement TM)

Capture the NH3 in clean water and you can get up to 25 % NH3 solution into the bargain!

The sodium nitrate can be recrystallised easily: sol. at 90 C is 148 g/100 ml, at 0 C it’s 73 g/100 ml, 50 % yield.

Any unreacted AN is highly soluble and stays in the mother liquor (which could be reused several times to minimise NaNO3 loss). Insolubles can be hot filtered off.

Oooops: looks like woelen beat me to it but with potash instead of caustic soda!

[Edited on 26-2-2013 by blogfast25]

AJKOER - 26-2-2013 at 16:02

OK, in my humble opinion, in suggesting a purification routine, perhaps one should always be very clear of what is the precise nature of the compound to be purified and its intended use. The purification approach should then follow based on meeting necessary standards to avoid critical impurities imperiling safety and in efficiently achieving a suitable product.

For example, if the intended use is for an energetic demonstration, this could be particularly problematic if unwanted metal impurities were present. This could lead to a particularly unstable and dangerous mixture.

Now in the washing/recrystallization synthesis suggested above, if an Fe impurity is present that is deemed problematic, then one or more additional steps may have to be implemented to remove the Iron. For example, an additional step could be to add a small amount of HNO3 (to remove soluble carbonates and sulfides) and then Barium iodide:

2 NH4NO3 + BaI2 --> Ba(NO3)2 (s) + 2 NH4I

and collect the relatively insoluble Barium nitrate and rinse. Then react with hot aqueous Na2CO3 to form NaNO3:

Na2CO3 + Ba(NO3)2 ---> 2 NaNO3 + BaCO3 (s)

If no similar solution possible, then perhaps a more drastic approach, as I detailed based on NO extraction, may be in order.

[Edited on 27-2-2013 by AJKOER]

ElectroWin - 26-2-2013 at 19:27

remember that pure ammonium nitrate is about 35% nitrogen.

DONALD W - 26-2-2013 at 19:35

I have try to recrystallize the 55 lb bags of Ammonia Nitrate fertilizer several times and finally after reading up on it--gave up. It appear they have chemically alter the fertilizer so that when you do recrystallize it recrystallize as a urea salt. Beautiful crystal but can not use to synthesize nitric acid. If you ever find out the secret, please let us all know

blogfast25 - 27-2-2013 at 05:48

Quote: Originally posted by AJKOER

OK, in my humble opinion, in suggesting a purification routine, perhaps one should always be very clear of what is the precise nature of the compound to be purified and its intended use. The purification approach should then follow based on meeting necessary standards to avoid critical impurities imperiling safety and in efficiently achieving a suitable product.

For example, if the intended use is for an energetic demonstration, this could be particularly problematic if unwanted metal impurities were present. This could lead to a particularly unstable and dangerous mixture.

Now in the washing/recrystallization synthesis suggested above, if an Fe impurity is present that is deemed problematic, then one or more additional steps may have to be implemented to remove the Iron. For example, an additional step could be to add a small amount of HNO3 (to remove soluble carbonates and sulfides) and then Barium iodide:

2 NH4NO3 + BaI2 --> Ba(NO3)2 (s) + 2 NH4I

and collect the relatively insoluble Barium nitrate and rinse. Then react with hot aqueous Na2CO3 to form NaNO3:

Na2CO3 + Ba(NO3)2 ---> 2 NaNO3 + BaCO3 (s)

If no similar solution possible, then perhaps a more drastic approach, as I detailed based on NO extraction, may be in order.

[Edited on 27-2-2013 by AJKOER]

Yeah, you're right, spend at ton on BaI2 for doing something that's otherwise straightforward. You sound like like the kind of guy would make a rectangle first to fabricate a wheel.

And what's this Fe thing about? Go on, indulge me and dig up some obscure stuff, just for laughs.

AJKOER - 27-2-2013 at 08:32

Blogfast:

You should know that metallic impurities in oxidizers could be a problem. I have previously documented on this forum of a cargo ship transporting a relatively safe hypochlorite, which have a small Mg impurities. The resulting self sustaining fire and large insurance loss were well investigated.

Now, not all impurities are equal so knowing how the product will be transported, stored and employed are important points.
--------------------------------------

Now on my suggested use of BaI2, which is one of the most soluble Barium salts. FYI, a Ba salt can be purchased from pottery supply stores relatively cheaply. NaI is available from eBay, so one could make BaI2, or a less soluble candidate, perhaps BaCl2, if the iodine is an issue. The concept of trying to precipitate the nitrate out of the solution is not perfect (that is why the solution should be pretreated with an acid and boiled to remove carbonates, etc..), but I do like the idea of seeing (via the precipitate) the potential maximum amount of nitrate available.

[Edited on 27-2-2013 by AJKOER]

woelen - 27-2-2013 at 09:23

@AJKOER: Please try to be more practical. I have the impression that nearly all of your chemistry is extremely impractical. I agree with blogfast25 that using BaI2 to get pure nitrate out of fertilizer grade ammonium nitrate is totally impractical. Never ever would I perform the process you mention.

The same is true for most of your other posts. Sometimes they really irk me. As a recent example comes to my mind the making of chlorates or chloric acid, involving distillation of HOCl. Immensely impractical and hardly possible to do in real life.

Have you done any practical chemistry, either as a hobbyist, or in your work? If you have, then you should know how impractical your suggestions are!

plante1999 - 27-2-2013 at 09:31

AJKOER post are impratical most of the time, and very often about a bottle of bleach and one of vinegar, too often. Like someone already said, everytime I look at a bottle of bleach, I can't think to AJKOER hypochloprite reaction. To be honest, I think he desrve a title about all these hypotetical bleach related posts.

AJKOER - 27-2-2013 at 10:21

Quote: Originally posted by DONALD W

I have try to recrystallize the 55 lb bags of Ammonia Nitrate fertilizer several times and finally after reading up on it--gave up. It appear they have chemically alter the fertilizer so that when you do recrystallize it recrystallize as a urea salt. Beautiful crystal but can not use to synthesize nitric acid. If you ever find out the secret, please let us all know

Here is an idea of you still have any of your altered (not surprised) NH4NO3 around. Take just a small amount of the dry altered NH4NO3, mixed well with Oxalic acid dihydrate (H2C2O4.2H2O) and heat. This is based on the reported reaction upon heating of dry mixed NaCl with H2C2O4 releasing CO and HCl vapors.

Hoped for results: vapors of HNO3, H2O, NO2, CO and CO2. Condense vapors into cold water containing Na2CO3 forming NaNO3. It is hoped that the H2C2O4 will decompose the unwanted organic additives.

Worst case: the mixture ignites/explode, which is more likely using anhydrous H2C2O4 (more sensitive and expensive) in a confined environment. Note, the otherwise stable H2C2O4 can be detonated and contribute to the explosion, and as such, scale is important. Also, a big downsize would be if the vapors included HCN which can be formed by the combustion of hydrocarbons in the presence of ammonia (from the thermal decomposition of NH4NO3). For example, with CH4 as an illustration:

2 CH4 + 2 NH3 + 3 O2 → 2 HCN + 6 H2O

By dissolving in aqueous Na2CO3, NaCN could be formed. To address this issue, one could replace the aqueous Na2CO3 with dilute HNO3. After gas collection, add BaI2 to create a precipitate of Ba(NO3)2, but the highly soluble Ba(CN)2 should remain in solution.

In any event, perform safely (outdoors) as the fumes are expected to be corrosive and toxic.

[Edited on 27-2-2013 by AJKOER]

blogfast25 - 27-2-2013 at 13:14

You're just taking the Micky, AJ. Please get lost.

garage chemist - 27-2-2013 at 13:32

Has someone already tried recrystallizing AN from methanol or ethanol? The solubility in these solvents is much more desirable than that in water.
Solubility in MeOH is 20% at 30°C and 40% at 60°C. I'd see this as a promising starting point.
In Ethanol the solubility is 4% at 20°C, with no information about the solubility at higher temperatures.

AJKOER - 27-2-2013 at 13:36

Well Blogfast, I could get lost if you would/could propose a real solution assuming the alteration is indeed correct.

To possibly restore your confidence, note this quote:

"When common salt is distilled with aqueous oxalic acid, a large quantity of hydrochloric acid is evolved. (Berthollet, Statique Chim. 1, 271.) Dry chloride of sodium or chloride of calcium intimately mixed with hydrated oxalic acid, gives off all its hydrochloric acid when heated, so that the residue left after ignition consists of carbonate of soda or carbonate of lime. (A. H. Wood, Phil. Mag. J. 5, 445; compare Kobell (J. pr. Chem. 14, 379.)"

Source: "Hand-book of Chemistry", Volume 9, by Leopold Gmelin, page 120. Link: http://books.google.com/books?pg=PA120&lpg=PA120&dq=...
-------------------------------------

Donald W, please verify that you purchased a product that contains NH4NO3 and not just Urea.

Next, assuming you have, here is another old method from a prior thread on this topic (see http://www.sciencemadness.org/talk/viewthread.php?tid=1112 ) that may still work. To quote from Stanfield:

"I've found a french text on google on how to purify it, here it is :

"dissolve the fertilizer in hot methanol and filter the solution. By mixing the solution with an equal volume of unleaded gasoline, the ammonium nitrate will instantly cristalize."

[EDIT]: Just noticed that Garage Chemist has just posted a similar idea (Blogfast, do you think he is on the micky too?)

[Edited on 27-2-2013 by AJKOER]

garage chemist - 27-2-2013 at 13:44

I do not see how adding gasoline to precipitate the AN would not precipitate the impurities as well. Recrystallization is generally done by lowering the temperature on a hot saturated solution, and only in very special cases by adding a nonsolvent.

DraconicAcid - 27-2-2013 at 13:57

Quote: Originally posted by garage chemist

I do not see how adding gasoline to precipitate the AN would not precipitate the impurities as well. Recrystallization is generally done by lowering the temperature on a hot saturated solution, and only in very special cases by adding a nonsolvent.

Depends on what the impurities are.

Remember, the impurities generally stay dissolved in a recrystallization, not because they are more soluble in the final solvent, but because they start out at a lower concentration.

garage chemist - 27-2-2013 at 14:18

If the impurity is calcium nitrate, adding gasoline will precipitate it as well. If it is to stay in solution, the amount of gasoline will have to be very carefully adjusted.

By mixing two solvents, you create a waste solution which is difficult to recycle.
If you recrystallize from a single solvent by temperature change, you can reclaim your solvent pure by distilling it off.

blogfast25 - 27-2-2013 at 14:34

Quote: Originally posted by AJKOER

To possibly restore your confidence, note this quote:

If anyone here has a confidence problem it's you. You know diddlysquat about chemistry and your sole purpose here is to bore people with your impossible schemes and obscure references. I doubt seriously if you own as much as a single test tube.

Do take up knitting.

AJKOER - 27-2-2013 at 16:37

Quote: Originally posted by blogfast25

Quote: Originally posted by AJKOER

To possibly restore your confidence, note this quote:

Upon more research, perhaps an easier and less expensive synthesis. First idea inspired by this from Science Labs' MSDS on BaCO3. To quote (link: http://www.sciencelab.com/msds.php?msdsId=9927090 ):

"Solubility [BaCO3]:
Very slightly soluble in cold water. Solubility in water: 0.024 g/l; 0.0022 g/l @ 18 deg. CAlmost insoluble in water. Soluble in solution of dilute hydrochloric acid, nitric acid, or acetic acid. Soluble in solution of ammonium chloride or ammonium nitrate. Insoluble in sulfuric acid. Soluble in ethanol"

So as BaCO3 (cost around $3.00 per lb from Pottery supply store) is soluble in NH4NO3, dissolve and remove any immediate precipitate (impurities). Reaction:

BaCO3 + 2 NH4NO3 <----> NH4(CO3)2 (aq) + Ba(NO3)2 (aq)

Upon boiling the solution (avoid fumes), Ammonium carbonate (or bicarbonate) decomposes moving the reaction to the right. Cool and add water and Barium nitrate could fall out of solution.
To purify further dissolve the Ba(NO3)2 in ethanol. React with Na2CO3 to form NaNO3 and regenerate the Barium carbonate.
----------------------------------------------

Also, someone may get ideas from this on NH4NO3 (linkhttp://www.sciencelab.com/msds.php?msdsId=9927336 ):

"Solubility[NH4NO3]:
Easily soluble in cold water, hot water. Soluble in acetone. Partially soluble in methanol. Insoluble in diethyl ether."
-----------------------------------------------

[EDIT] Now Blogfast, per my posted thread on Ferrates where I displayed my fine Italian glassware (see it again at http://www.sciencemadness.org/talk/viewthread.php?tid=17280#... ), your comments are as accurate as usual.

[Edited on 28-2-2013 by AJKOER]

ScienceSquirrel - 27-2-2013 at 17:20

Quote: Originally posted by AJKOER

Quote: Originally posted by DONALD W

The reaction of ammonia, methane and oxygen to form hydrogen cyanide will only take place in the presence of a platinum catalyst.
The uncatalysed reaction will form water, nitrogen and carbon dioxide and monoxide depending on the stochiometry.
Hydrogen cyanide will not be a significant product.

dulio - 27-2-2013 at 18:19

Well, thank you all folks. Things are getting clearer. By now, I would address AJKOER first. I would prefer wolen's method. Keep I mind the nitrate will be used for nitric acid production.

AJKOER - 28-2-2013 at 07:06

Quote: Originally posted by ScienceSquirrel

Yes, I agree on the need for the catalyst for CH4 which I cited as an illustration of a possible similar chemical reaction that could form HCN as, for example, commonly occurs with the combustion of plastics. The nature of the presumed organic (?) additive in the current case is apparently unknown.

However, the real unsaid reason I would not rule out the formation of HCN, as we could be addressing a serious effort to discourage the recovery of NH4NO3, is a conceivably conscious intention to allow the possible formation of HCN (during a thermal approach) should not be ruled out. In the same vain, my initial though was that a serious complex effort was the only avenue to unlock the NH4NO3 (hence the difficult, dangerous and costly thermal approach with Oxalic acid and BaI2, although the aqueous H2C2O4/Cu route is perhaps less dangerous and costly, but a more prolonged procedure).

Now, for all my critics, my simple suggested use of BaCO3 with boiling and dilution to recover Ba(NO3)2, as a possible path to another nitrate or HNO3, is not overly complex or costly, in my opinion. However, Barium salts are toxic and require care in handling. But, in comparison to working with Methanol and gasoline, perhaps not much difference assuming the latter is at all successful as well.
-----------------------------------------------------

Now, those who are irked that I, in a prior thread, cited that HOCl can indeed be used to form HClO3 should re-read the previously supplied World Patent employed in an actual large scale commercial production of Chloric acid. The key is chloride free HOCl in a closed system. The fact that AgOCl readily forms AgClO3 and an insoluble (in effect, removed) AgCl is an instructive comparison as to how the reaction can proceed.

[Edited on 28-2-2013 by AJKOER]

ScienceSquirrel - 28-2-2013 at 07:30

Quote: Originally posted by AJKOER

Yes, I agree on the need for the catalyst for CH4 which I cited as an illustration of a possible similar chemical reaction that could form HCN as, for example, commonly occurs with the combustion of plastics.

However, the real unsaid reason I would not rule out the formation of HCN, as we could be addressing a serious effort to discourage the recovery of NH4NO3, is a conceivably conscious intention to allow the possible formation of HCN (during a thermal approach) should not be ruled out. In the same vain, my initial though was that a serious complex effort was the only avenue to unlock the NH4NO3 (hence the difficult, dangerous and costly thermal approach with Oxalic acid and BaI2).

Plastics that form hydrogen cyanide on combustion are typically polymers of acrylonitrile. Hydrogen cyanide is eliminated from the backbone of the polymer as the plastic pyrolyses. Reaction with nitrogen from the air does not take place.
Gaseous nitrogen is inert to almost all reactants apart from alkali metals, a few transition metal complexes, etc. That is why it is used as a blanketing gas in air sensitive chemistry when the more expensive argon is not justified.
If it was not so unreactive millions of pounds and thousands of hours of research time would not have been expended on the problem of nitrogen fixation.
If someone has crude ammonium nitrate and they want to make a pure nitrate salt from it, the easy way is to add potassium carbonate or hydroxide, boil to drive off the ammonia and then crystallise.
Good old fashioned nonhygroscopic potassium nitrate, usable in all sorts of things.
Easy peasy, lemon squeezy and no arsing around with large quantities of toxic barium salts.

AJKOER - 28-2-2013 at 10:04

Quote: Originally posted by ScienceSquirrel

....
If someone has crude ammonium nitrate and they want to make a pure nitrate salt from it, the easy way is to add potassium carbonate or hydroxide, boil to drive off the ammonia and then crystallise.
Good old fashioned nonhygroscopic potassium nitrate, usable in all sorts of things.
Easy peasy, lemon squeezy and no arsing around with large quantities of toxic barium salts.

Perhaps a closer examination of the K2CO3/KOH approach is in order:

1. Safety: Even caustic KOH is preferable over a dissolved Barium salt (very toxic). The dry BaCO3 is a little better due to its low solubility, but my preference would still for KOH/K2CO3 with respect to safety.

2. Cost: K2CO3 purchased in bulk (50 lbs) is as about the same price as BaCO3, otherwise a little more expensive. However, there is an underlying assumption of equal access to both products, which may be a function geography.

3. Chemistry: is basically the same with the formation of Ammonium carbonate (or bicarbonate) which decomposes on boiling moving the reaction to the right.

4. Impurities: The immediate presence of a dissolved Barium salt could permit the immediate removal of, say sulfates, as a precipitate, or one of a wide range of other possibilities. Not the case for KOH/K2CO3. Also, Ba(NO3)2 is soluble in ethanol while KNO3 is only slightly soluble in ethanol, which could permit the formation of a purer product. The Barium approach appears a little better here.

5. Yield: The expected yield is most likely related to solubility differences between Ba(NO3)2 ( which is 6.77 g/100g at 10 C) and KNO3 ( 47.00 g/100g at 10 C). Here the Barium approach appears better.

So, other than safety (certainty important and over-riding in some circumstances), I think the edge goes to the Barium carbonate approach.

[Edited on 28-2-2013 by AJKOER]

blogfast25 - 28-2-2013 at 14:13

Quote: Originally posted by AJKOER

So, other than safety (certainty important and over-riding in some circumstances), I think the edge goes to the Barium carbonate approach.

[Edited on 28-2-2013 by AJKOER]

Why, what's the worst that could happen in your case? A paper cut? Getting your finger trapped in your laptop?

AJKOER - 28-2-2013 at 14:24

Actually Blogfast, in some states one can buy NH4NO3 exploding targets, or just the NH4NO3. You should see some of the videos, cool!

No need to get me or my expensive glassware contaminated, just pure fun

-------------------------------------

I noticed you did not comment on the relative merits of the BaCO3 versus K2CO3 proposals, interesting.

By the way, it occurred to me to make things even easier, one does not have to even boil! After dissolving the BaCO3 with the addition of hot water, remove any impurities (precipitates) and let the solution sit in an wide mouth open vessel in the sun. Ammonium carbonate may decompose by itself with evaporation. Then, the Ba(NO3)2 may precipitate by itself or upon periodic additions of water.

But if one really like to inhale NH3 fumes, then boil away to concentrate in the hope of obtaining a KNO3 precipitate after cooling (don't forget the ice). Somehow, the K2CO3 procedure doesn't sound too cool.

[Edited on 28-2-2013 by AJKOER]

blogfast25 - 1-3-2013 at 05:20

We all know about the potentially explosive nature of ammonium nitrate. Yet no one sane of mind would use BaCO3 as an alkali instead of KOH for the alkali displacement of AN. You would. The target here is NaNO3, not NH4NO3.

AJKOER - 1-3-2013 at 06:29

Blogfast:

The answer to my question as to the author's goal was a nitrate (NaNO3) for the production of HNO3.

So my original response of treatment with aqueous H2C2O4 followed by Cu to NO to NO2 to HNO3 is still a good idea for the creation of pure conc Nitric acid, especially if altering agents have been introduced (as Oxalic acid acts as a reducing agent on many organic compounds). The process, however, requires some degree of expertise owning to the collection and employment of corrosive and toxic gases.

Now, Barium nitrate to HNO3, again the application of a dilute solution of H2C2O4 and filter. Why dilute? Read my thread on the use of H2C2O4 to produce strong acids and some of the adverse safety issues that were observed. So my personal recommendation on the use of H2C2O4 to form conc acids is to target more dilute versions, not because the procedure does not work, but because, at times, it works too well and many are not ready for, or have knowledge of, or experience with (including seasoned chemists) the associated bad behavior of many very conc acids (as they are just not available).

[Edited on 1-3-2013 by AJKOER]

blogfast25 - 1-3-2013 at 07:39

Quote: Originally posted by dulio

Anyway, let us start from the basics. How can I purify the nitrate in order to further convert it into sodium nitrate? I also want to keep some pure stuff in my shelf.

Any help is welcome. Thanks in advance, people.

It doesn't get much clearer than that.

[Edited on 1-3-2013 by blogfast25]

AJKOER - 1-3-2013 at 08:20

Quote: Originally posted by dulio

Well, thank you all folks. Things are getting clearer. By now, I would address AJKOER first. I would prefer wolen's method. Keep I mind the nitrate will be used for nitric acid production.

No Blogfast, it doesn't get any clearer than this.

ScienceSquirrel - 1-3-2013 at 10:03

Barium carbonate is only slightly soluble in water and is only feebly basic, it will not displace ammonia from it's salts.
You can boil ammonium nitrate solution with barium carbonate until the cows come home, no reaction will happen.

blogfast25 - 1-3-2013 at 11:14

Quote: Originally posted by ScienceSquirrel

Barium carbonate is only slightly soluble in water and is only feebly basic, it will not displace ammonia from it's salts.
You can boil ammonium nitrate solution with barium carbonate until the cows come home, no reaction will happen.

Not even if you throw a belief system at it?

AJKOER - 1-3-2013 at 15:03

OK, is anybody acquainted with the use of NH4Cl/NH4NO3 to help in commercial processes dealing with the dissolving of Ca(OH)2?

The explanation of why nearly insoluble salt increase their solubility in the present of strongly ionic salts is complex. One theory cites the increase in ionic interactions from the "non-common ion effect". As a result, "a sparingly-soluble salt will be more soluble in a solution that contains non-participating ions." per this educational reference, http://www.chem1.com/acad/webtext/solut/solut-6b.html .

This is also referred to as the "Salt Effect" to quote from Wikipedia:

"The salt effect[2] refers to the fact that the presence of a salt which has no ion in common with the solute, has an effect on the ionic strength of the solution and hence on activity coefficients, so that the equilibrium constant, expressed as a concentration quotient, changes". Please see Wikipedia comments on solubility equilibrium at http://en.wikipedia.org/wiki/Solubility_equilibrium ).

For more advanced details see http://www.jim.or.jp/journal/e/pdf3/45/04/1317.pdf . To quote from the abstract:

"We developed a chemical model to analyze ionic equilibria in a cobalt chloride solution at 298K. The chemical model consisted of chemical equilibria, mass and charge balance equations. The activity coefficients of solutes and water activity were calculated with Bromley equation. Values of the equilibrium constants for the formation of cobalt chloride complexes at zero ionic strength and of the interaction parameters were estimated by applying Bromley equation to the reported equilibrium constants at different ionic strength". Now, in the current context, note that NH4Cl/NO3 are salts of a weak base and strong acid, which are highly ionic and have correspondingly low pHs."

This topic actually can up on another forum (please see http://www.chemicalforums.com/index.php?topic=65345.msg23855... )

Now the solubility of BaCO3 in NH4Cl/NO3 is apparently (or, at least cited to) increased per a ScienceLab MSDS on Barium carbonate as I previously noted. This theoretically parallels the case of the increased solubility of Ca(OH)2 in NH4Cl/NO3 discussed above. So does anybody have a source to indicate this is, in fact, not correct?

Note, I do not dispute the observation/fact that both BaCO3 and Ca(OH)2 are otherwise very insoluble in water.

[Edited on 2-3-2013 by AJKOER]

ScienceSquirrel - 1-3-2013 at 15:50

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.
That makes it at least 70 times less soluble and 1,000 times weaker as a base.
Do some experiments.
Boil calcium hydroxide with an ammonium salt and smell the ammonia and then boil calcium or barium carbonate with an ammonium salt and smell arse all.
Humanity faces huge challenges; the rising cost and availability of fuels, possible ciimate change and the availability of food and water.
The world may need good science if we are to survive as a species.

DraconicAcid - 1-3-2013 at 16:08

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.)

ScienceSquirrel - 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.

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.

AJKOER - 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]

dulio - 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.

Fyndium - 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]

Fyndium - 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.