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dulio
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Ammonium nitrate purification
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.
\"Do not compare you to others, but with the best which yourself can do.\"
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elementcollector1
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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.
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cyanureeves
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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.
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jock88
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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.
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BromicAcid
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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.
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dulio
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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 PO43-.
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.
\"Do not compare you to others, but with the best which yourself can do.\"
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dulio
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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.
\"Do not compare you to others, but with the best which yourself can do.\"
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dulio
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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.
\"Do not compare you to others, but with the best which yourself can do.\"
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BromicAcid
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@ 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.
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dulio
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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.
\"Do not compare you to others, but with the best which yourself can do.\"
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blogfast25
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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]
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elementcollector1
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Still, chilling a saturated solution of AN at 100 C to 0 C causes about 9/10 of the AN to precipitate out...
Elements Collected:52/87
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blogfast25
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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]
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AJKOER
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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.
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weiming1998
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| 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]
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AJKOER
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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).
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woelen
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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]
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blogfast25
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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]
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AJKOER
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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]
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ElectroWin
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remember that pure ammonium nitrate is about 35% nitrogen.
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DONALD W
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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
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blogfast25
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| 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.
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AJKOER
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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]
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woelen
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@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!
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plante1999
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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.
I never asked for this.
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