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AJKOER
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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]
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blogfast25
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You're just taking the Micky, AJ. Please get lost.
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garage chemist
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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.
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AJKOER
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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]
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garage chemist
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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.
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DraconicAcid
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| 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.
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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.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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garage chemist
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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.
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blogfast25
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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.
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AJKOER
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| Quote: Originally posted by blogfast25 |
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. |
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]
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ScienceSquirrel
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| Quote: Originally posted by AJKOER | | 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] |
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.
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dulio
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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.
\"Do not compare you to others, but with the best which yourself can do.\"
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AJKOER
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| Quote: Originally posted by ScienceSquirrel | 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. |
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]
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ScienceSquirrel
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| 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).
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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.
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AJKOER
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| 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]
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blogfast25
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| 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?
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AJKOER
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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]
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blogfast25
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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.
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AJKOER
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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]
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blogfast25
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| 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]
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AJKOER
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| 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.
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ScienceSquirrel
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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.
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blogfast25
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| 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? 
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AJKOER
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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]
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ScienceSquirrel
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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.
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DraconicAcid
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| 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.)
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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