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mbrown3391
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Ammonium nitrate without nitric acid
Supposed I dissolve some potassium nitrate into some store-bought ammonium hydroxide solution and then slowly drip sulfuric acid into this solution.
Would a significant amount of ammonium nitrate be formed? Unfortunately i do not have access to nitric acid so im looking for a way to make ammonium
nitrate without it. Any other ideas would be appreciated.
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garage chemist
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Ammonium nitrate is way more soluble than potassium nitrate/sulfate, so you wont have any luck with that. Why do you want to make ammonium nitrate
from KNO3? I always did it the other way around, with AN (isolated from AN/chalk fertilizer) which was boiled with K2CO3, giving off NH3 and CO2 and
leaving behind a KNO3 solution.
But as you have sulfuric acid and KNO3, you essentially have nitric acid. You just need a distillation setup.
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YT2095
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not taking part in this.
[Edited on 20-6-2007 by YT2095]
\"In a world full of wonders mankind has managed to invent boredom\" - Death
Twinkies don\'t have a shelf life. They have a half-life! -Caine (a friend of mine)
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mbrown3391
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I would distill the nitric acid, but i dont have an all glass distillation apparatus. How would i make ammonium nitrate with ammonium chloride?
Edit: I definitely don't need to make more KNO3. I have over 10 pounds of it originally purchased for gunpowder. I guess i over ordered!
[Edited on 20-6-2007 by mbrown3391]
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YT2095
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not taking part in this.
[Edited on 20-6-2007 by YT2095]
\"In a world full of wonders mankind has managed to invent boredom\" - Death
Twinkies don\'t have a shelf life. They have a half-life! -Caine (a friend of mine)
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mbrown3391
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what do you mean. Why aren't you taking part in this?
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Aqua_Fortis_100%
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Alternativelly to distillation method, you can try the
US patent 3981975 : http://www.freepatentsonline.com/3981975.html
by doing this you probably will dont need any expense apparatus and the danger of being poisoned by nasty NOx fumes and such..
the only things you will need are dichloromethane (DCM) , xNO3-H2SO4(in your case the KNO3) , a glass funnel , some glass wool and dry ice..
simple extract the anhydrous HNO3 (or almost anhydrous..(DCM has an azeotrope with water..but you can simple overcome this by using any suitable
dehydrating agent with the DCM))with DCM from saltpeter-sulfuric acid mix , pour off the DCM-HNO3 mix, and chill in a separate container with some dry
ice sorrounding the container.. crystalline anhydrous HNO3 will precipitate and you can then filtrate it in your previous chilled glass funnel-glass
wool .. after you could just let react CAREFULLY with ammonia (i think which this will generate LOTS of NH4NO3 smoke, so do this outside) and then
you will get your ammonium nitrate, pure and nice 
(although i think which you can make far more great things with anhydrous HNO3 than ammonium nitrate...)
the best thing of this is which you can just recycle the DCM by simple distillation in very improvised (? ) destiller apparatus..
this is one of my future projects ..unfortunatelly i dont know ANY OTC solvent where i live which contains DCM..the label of almost ALL products are
uninformational of the contents..(usually they say "product made from petroleum hydrocarbons(...)"  )
for sources, you can also see this:
http://www.roguesci.org/theforum/showthread.php?t=1910
and, of course, search on this board (SMDB) to get more info about..
any additional ideas ?
[Edited on 20-6-2007 by Aqua_Fortis_100%]
[Edited on 20-6-2007 by Aqua_Fortis_100%]
"The secret of freedom lies in educating people, whereas the secret of tyranny is in keeping them ignorant."
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YT2095
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if you have 10 pounds LEFT OVER! from making BP, and then want to make AN but ask such basic questions, that can only mean bad things 
\"In a world full of wonders mankind has managed to invent boredom\" - Death
Twinkies don\'t have a shelf life. They have a half-life! -Caine (a friend of mine)
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mbrown3391
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im not using the ammonium nitrate as an explosive if thats what your thinking. Just because i have 10 pounds doesn't mean im going to use it all for
this
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not_important
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Go to garden supply shops, look for a bag of ammonium sulfate fertilizer. Make saturated solutions of (NH4)2SO4 and KNO3, filter the solutions hot,
then measure and mix in the proper proportions. The best way to do this is use a large waterbath just below boiling, keep all the containers hot, even
do the filtering with the receiving flask in the bath and the funnel pre-heated in the bath and warmed by the steam coming off.
K2SO4 is much less soluble than (NH4)2SO4 or either of the nitrates, some of it will precipitate out. Turn off the heat, allow the solution to cool
slowly - you can leave it in the bath. If you take it out you'll have to keep swirling or stirring it. At the proper temperature (*) decant the
solution from the K2SO4 through a filter, again heated to at least the temperature of the solution. Let cool further.
You will have to look up the solubilities on your own, convert to molar solubilities, and draw the curves. Using saturated solutions means you don't
need to weigh the reactants, just measure volumes of solutions. Remember that when you mix them you're diluting them as well - you have to total
the amount of water and calculate the concentrations in the mixed solutions.
(*) As I remember it, K2SO4 is the least soluble of the four salts all the way down to zero, and after adjusting concentrations none of the other
salts will reach saturation, so maybe you can skip filtering part way down the cooling curve. If that's is the case, just cool and decant/filter, if
not you'll want to filter before you reach saturation of any of the other salts.
You'll end up evaporating the solution to get the NH4NO3 to crystallise out. Don't try to be greedy, leave some mother liquor to hold impurities; and
most certainly don't "boil it down", at the most use one of those warmer trays that don't get hotter than 40 C or so.
Recrystallise by making a hot saturated solution and letting it cool. Alternatively dissolve the NH4NO3 in boiling methanol and filter. it's
moderately soluble in MeOH while KNO3 is only sparingly so and the sulfates are nearly insoluble. Again, you'll have to research the precise values
of the different solubilities yourself so you use just enough alcohol to dissolve the NH4NO3 and leave the other stuff behind.
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mbrown3391
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thanks. that sounds like the best option so far.
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Aqua_Fortis_100%
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oh... thanks also..this method using ammonium sulphate is one more alternative to try..
but best than using KNO3 is using calcium nitrate , Ca(NO3)2 ,and as ammonium sulphate, is also readily avaliable in some garden centers and very
cheap (although here i still dont seem any of this in the shelves of garden houses ..)
make hot saturated solutions of both and mix then..most of the calcium sulphate ppt readily:
Ca(NO3)2 + (NH4)2SO4 ---> CaSO4 + 2 NH4NO3
filter , and then you can let cool to room temperature and chill more in a freeze.. filter again and boil of some water ( e.g. half of initial..i'm
not sure right now what volume is best), let cool to room temperature and filter again.. boil some more until almost of the liquid has boiled off
and chill to lowest temperature you can get to obtain more ammonium nitrate.. filter , dry this is and hotplate and store airtight...(keep in mind
which your ammonium nitrate has some impurities on it... this can be dangerous in making explosives because the sensibility of ammonium nitrate
increases greatly when impurities (even in order of 0.5 %, etc) is present..so beware and search more on this board to get the right words on the
matter..
(another way would be mix saturated solution of Ca(NO3)2 in dilute sulfuric acid (filtrated battery acid) and filter the new CaSO4 formed in an glass
funnel with some glass wool in it (because of the nitric acid formed) and react with ammonia in excess and get rid of water, just boiling to half and
chill down, filter in common paper filter the newest CaSO4 and boil more and chill and then filtering the ammonium nitrate)
"The secret of freedom lies in educating people, whereas the secret of tyranny is in keeping them ignorant."
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not_important
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The problem with the calcium nitrate/ammonium sulfate route, having done this when I was 12, is that calcium sulfate has its solubility increased by
ammonium salts. If you take ammonium nitrate made this way, and recrystallised several times, it still leaves a fair residue of calcium oxide behind
when you heat a bit to decomposition. Depends on your application if the calcium is a problem or not.
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chemkid
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Lots of cold packs
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jimmyboy
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not_important --- how would you get the ammo nitrate to "xtallize out?" -- wouldn't potassium sulfate be the first to crash at saturation?
When I tried the kno3/nh4so4 route years ago I added alcohol thinking i would get ammo nitrate left in the solution -- all that happened was kno3
crashed out in tiny crystals - this reaction is useless from what I can tell because the solubilities of k2so4 and kno3 are too close - you really
need an insoluble product to get things going (Ca sulfate)
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not_important
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| Quote: | Originally posted by jimmyboy
not_important --- how would you get the ammo nitrate to "xtallize out?" -- wouldn't potassium sulfate be the first to crash at saturation? ...
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You need to getthe solution concentrated enough that K2SO4 starts to crystallise out, in a hot solution it has by far the lowest solubilities of all
four salts. The idea is to adjust thing so that the solution ends up almost saturated in one of the other salts, with much of the K+ and SO4(-2) as a
solid. Remove the K2SO4, continue to evaporate and/or cool.
You need to draw the solubilities curves, doing so in terms of moles of each salt so you can see the relationships better. Sometimes in cases like
this you'll plot the solubility of "K2(NO3)2" in order to make it easier to see things - moles of (NH4)2SO4 and (KNO3)2 vs (NH4NO3)2 and K2SO4.
It isn't always A + B => C + D(xtals). When NaCl is a product you may take advantage of its fair flat solubility curve - mixing the two reagents
near boiling and getting a precipitate of NaCl, separating that, then cooling to get the other product to crystallise out with only a small amount of
NaCl contaminating it. And yes that means you need to recrystallise the product to get rid of the NaCl that does come down with it.
At 100 C, 100 ml water will dissolve roughly 11 moles of NH4NO3, 2 moles of KNO3, 3/4 of a mole of (NH4)2SO4, or 1/8 mole K2SO4. If you mix 100 ml of
saturated (NH4)2SO4 and 75 ml of saturated KNO3 solutions @ 100 C, then about 1/2 of a mole of K2SO4 will drop out.
You now have only 1/4 mole of SO4(-2) and 1/2 mole of K(+1) left. You could then evaporate, @ 100 C, the solution down until nearing the saturation
point for (NH4)2SO4, and again remove the solid K2SO4. Note that after the 1st removal of K2SO4 you can have not more than 1/4 mole of (NH4)2SO4 and
1/2 mole KNO3, both of which would dissolve in say 40 ml of water @ 100 C, while that 40 ml will hold less than 1/16 of a mole of K2SO4.
Continue to evaporate down to 25 ml, and the solution can hold no more than 3 grams of K2SO4. It would also be over-saturated in KNO3 and (NH4)2SO4
had you not been removing K2SO4 along the way; because you did so there's too little SO4(-2) and K(+1) left for either of those other salts to
crystallise out.
If you cooled the solution to zero C, about 7/8 of the ammonium nitrate would crystallise out, giving you a mush that would be hard to clean, along
with a gram or 2 of K2SO4. It's better to cool part way, separate the crystals, cool further and get the next batch, and so on.
All of these calculations were done in my sleepy head, and not using my toes, so the numbers may be off. But the idea is there, do the calculations
yourself and plot them.
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DeAdFX
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| Quote: | Originally posted by not_important
The problem with the calcium nitrate/ammonium sulfate route, having done this when I was 12, is that calcium sulfate has its solubility increased by
ammonium salts. If you take ammonium nitrate made this way, and recrystallised several times, it still leaves a fair residue of calcium oxide behind
when you heat a bit to decomposition. Depends on your application if the calcium is a problem or not. |
Does Calcium Phosphate suffer the same downfalls? Triammonium Phosphate should be a fairly easy to find at a large farming supply store correct?
I also wonder if its possible to make ammonium nitrate/nitric acid using mono or dihydrogen phosphates as can be done with bisulfate. This is just in
case there is no Triammonium phosphate.... I know that the hydrogens on the phosphate are of going to be of varying strength but am uncertain of how
it would apply to the problem...
Bisulfate is a weaker acid than nitric acid when I look at the acid strength charts. I guess the reaction proceeds because a precipitate is formed?
[Edited on 13-7-2007 by DeAdFX]
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jimmyboy
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Bisulfate will work for making the acid if heated dry -- but the topic is NOT nitric acid .. displacement reactions need one of three things to go
forward - an insoluble precipitate - a gas is formed or water is a product...
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not_important
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Don't know for the ortho phosphates.
While there is an triammonium orthophosphate, it loses ammonia quickly. The diammonium salt is easy to get, even it loses ammonia fairly readily.
Fertilizer grade ammonium phosphates are often not very pure, fluoride is a frequent contaminate. In that case, you might do better treating the
ammonium phosphate with sodium carbonate and a small amount of water, heating (~80C) to drive off NH2 and CO2, then leading those gases into a strong
solution of calcium nitrate. It's OK to have an excess of the NH3 and CO2, gently boiling the solution will remove the excess.
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len1
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Ammonium sulphate of the garden fertilizer variety is not very good for making other ammonium salts because it contains a large percentage of iron
II/III sulphate impurity. This is quite hard to get rid of.
The reason why CaSO4 dissolves in the presence of ammonium salts is that they are acidic and shift the HSO4 <-> SO4 equilibrium towards the
bisulphate. Calcium bisulphate is more soluble than CaSO4. In this regard a precipitate of the latter will easily dissolve when H2SO4 is added. In
the case of NH4+ ions in solution the ammonium salt of Ca(HSO4)2 is soluble.
Perhaps the way to shift the equilibrium back towards the much less soluble CaSO4 is to make the solution basic, say by adding NH3, and then filtering
the basic solution and evaporating off the NH3 and H20.
[Edited on 3-8-2007 by len1]
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len1
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I have now a way to purify ammonium sulphate of the garden fertilizer to a nice clean white powder.
The crude product is dissolved in H2O about 2.5kg to 6 liters, distilled water is best if you dont want to introduce Ca impurities. The solution is
filetered thru coarse tissue to remove any insoluble sulphates and organic charred mater (due to the H2SO4 used in its preparation), dont waste nice
filter paper on this filtration. The solution is now brown coloured due to Fe3+ present. Next add NH3 solution so the pH is about 8-9 (50ml 5% NH3
to 2l). There results after a few minutes a red-brown precipitate due to iron impurities present. The Fe2+ however remains in solutiuon to a
sufficient degree at this pH, reason being that it is relatively more soluble (sol. prod. 10^-16 for Fe(OH)2 ) compared to Fe(OH)3, which is one of
the most insoluble things known (10^-40). Next air is bubbled thru the solution from a compressed air source at a slow rate (a bubble a second) for a
few hours. This turns most of the Fe2+ to Fe3+, the euilibrium being shifted right due to the high pH. The solution is let settle until clear (2-3
days). It can not be filtered since the Fe(OH)3 is very fine and will pass thru almost any filter. Instead it is decanted by suction into am empty
bottle, leaving a small (1cm) layer at the bottom with the percipitate. The resulting solution is evaporated to give clean white (NH4)2SO4.
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not_important
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Solubility of CaSO4 in solutions of NH4NO3
grams/per liter 25 C
NH4NO3 CaSO4
0 ...... ...... 2,08
20 ...... ...... 3,7
40 ...... ...... 5,1
60 ...... ...... 6,0
80 ...... ...... 7,0
100 ...... ...... 7,65
150 ...... ...... 8,9
200 ...... ...... 9,85
300 ...... ...... 10,8
400 ...... ...... 11,4
600 ...... ...... 12,1
NH4Cl has a somewhat larger effect, perhaps 10 to 20 percent increase over the nitrate. At higher concentrations of ammonium salts the solubility of
CaSO4 begins to decline again.
Len, I did something dimly similar. I placed a large jar inside a fish tank, put a large test tube in the center of the jar and clamped it there,
filled the jar with fertiliser grade ammonium sulfate, filled the test tube with mineral oil and dropped in a cartridge header, and poured a cm or so
of fine sand on top of the sulfate. I then filled the tank with DW. The water was higher that the top of the jar, but the test tube stuck out of the
water. Finally I dropped in a bubble stone hooked up to an aquarium air pump, and added aqueous ammonia until pH 8.
I then applied power to the cartridge heater, slowly, until the oil reach 80 C. I then covered the tank and left it.
The warmed water inside the jar dissolved some of the sulphate. Eventual diffusion lead to the tank being filled with a saturated solution of the
sulfate, and then to crystals being deposited on the tank walls, cooler than the heater containing jar.
Eventually most of the sulfate in the jar dissolved. I turned off the heater, lifted the jar out of the tank and syringed out most of the solution in
the jar. The added fresh crude sulfate to the jar, and capped with fine sand again. Back in the tank, add aqueous ammonia again, turn the heater back
on.
I ran three 20 Kg sacks of ammonium sulfate through, got about 52 kg of recrystallised white (NH4)2SO4 off the tank walls and liters of the saturated
solution that I used in the garden for the next several years.
The cold/hot solubility ratio of ammonium sulfate was low enough that I'd gotten frustrated with typical recrystallisation. This way I lost less to
the mother liquor, and didn't have to put uf with filtering the solution.
[Edited on 13-8-2007 by not_important]
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len1
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Hi not important, very similar procedure indeed, and more imaginative. I wonder if carrying it out at 80C shifts the equilibrium any, and in what
way. How long does the diffusion take, and does the sand stop penetration of all insolubles.
On the subject of the solubility of Ca in the presence of NH4, I have a university textbook which states that this can also be achieved by addition of
H2SO4, and is due to the formation of the 'acid' H2Ca(SO4)2, in the first case, or its ammonium salt in the second. This is just calcium bisulphate,
which is more soluble than the sulphate. And since all ammonium salts with strong acids are acidic, I believe the dissolution is just due to the HSO4
<-> SO4 equilibrium being shifted left. In that case adding NH3 for basicity should reprecipitate the CaSO4 and not contaminate the salt any
Another possibility is to add a small amount of ammonium carbonate, sufficient to precipitate out the residual calcium ions as CaCO3. The pH needs to
be adjusted though so theres a sufficient concetration of carbonate and not bicarbonate
[Edited on 14-8-2007 by len1]
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not_important
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The solubilities in NH4Cl and NH4NO3 begin to decrease as the amounts of those salts increase above those in that table. And back when I was making
NH4NO3 that route, adding strong ammonia didn't seem to make much difference on the amount of calcium showing up, but I had no real way to get good
measurements of the amount of calcium in the ammonium nitrate - just eyeballing the residue after heating some on a steel sheet.
You might be able to add a small amount of urea, and then bring the solution to 90 Cl and hold it there for awhile, finally bringing it to a boil.
That would give both CO2 and NH3, boiling drives off excess.
I was letting things diffuse quite slowly, many days to process a several kg load. As there was very little current, the bubbler being located against
the tank wall as far from the jar as it could be, so the sand was a good barrier against insolubles drifting about. The aqueous sulfate will lose
ammonia at 80 C, but that was just the temperature in the inner container. By covering the tank the ammonia lose was kept down, I just added a bit of
NH3 aq each day as makeup.
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len1
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Not important, you are as always very useful, that method you suggested seems a jem, plus the table is useful. I wouldnt read too much into the
solubility of CaSO4 starting to drop at very large ammonium ion concentrations since judging from the table you gave us thats of the order of 6 mol/l
where theres substantial ion-ion interaction and simple thermodynamics fails. Adding NH3 not having any effect is another matter and would defeat the
mechanism for enhancing solubility I was suggesting. Since even small amounts of NH4 have an effect of CaSO4 solubility I doubt the effect is an
ion-ion interaction, plus Im sure Ca2+ does not form a coordination complex with ammonia. Certainly the content of the book I quoted does not explain
the reason for whats happening. I guess Ill have to try it out and see.
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