Sciencemadness Discussion Board - Ammonium nitrate without nitric acid

Ammonium nitrate without nitric acid

mbrown3391 - 20-6-2007 at 09:32

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.

garage chemist - 20-6-2007 at 09:36

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.

YT2095 - 20-6-2007 at 09:43

not taking part in this.

[Edited on 20-6-2007 by YT2095]

mbrown3391 - 20-6-2007 at 09:46

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]

YT2095 - 20-6-2007 at 10:21

not taking part in this.

[Edited on 20-6-2007 by YT2095]

mbrown3391 - 20-6-2007 at 10:30

what do you mean. Why aren't you taking part in this?

Aqua_Fortis_100% - 20-6-2007 at 10:55

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(...)" :mad:

)

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%]

YT2095 - 20-6-2007 at 10:56

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

mbrown3391 - 20-6-2007 at 11:22

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

not_important - 20-6-2007 at 11:52

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.

mbrown3391 - 20-6-2007 at 12:03

thanks. that sounds like the best option so far.

Aqua_Fortis_100% - 21-6-2007 at 04:16

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 :mad:

..)

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)

not_important - 21-6-2007 at 06:01

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.

chemkid - 23-6-2007 at 16:49

Lots of cold packs

jimmyboy - 11-7-2007 at 09:16

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)

not_important - 11-7-2007 at 10:27

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? ...

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.

DeAdFX - 13-7-2007 at 08:12

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]

jimmyboy - 13-7-2007 at 17:09

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...

not_important - 13-7-2007 at 18:33

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.

len1 - 1-8-2007 at 20:35

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]

len1 - 12-8-2007 at 20:57

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.

not_important - 12-8-2007 at 23:20

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]

len1 - 13-8-2007 at 23:22

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]

not_important - 14-8-2007 at 06:30

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.

len1 - 14-8-2007 at 17:25

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.

497 - 3-1-2009 at 20:06

I have a bunch of ~70% ammonium nitrate ~30% ammonium sulfate fertilizer that I am hoping to separate into relatively pure NH4NO3 and (NH4)2SO4. I'm wondering if it would be possible to use a system similar to the one not_important uses in the above post?

I was also wondering is I could add something else to the water to increase the solubility difference between the two salts. I'm thinking an alcohol or polyol might be helpful.. According to this ammonium sulfate's solubility is substantially reduced by the presence of various polyols. No actual numbers though. So this brings me to the question, does anyone have some actual data on water, alcohols or polyols and NH4NO3 and/or (NH4)2SO4 systems?

Basically I'm trying to find the optimum mixture that gives the largest solubility difference while still being cost effective and practical. Neat methanol seems like a possibility but has disadvantages like flamability (especially when working with large amounts

), cost/availability, and NH4NO3 not being very soluble in it. So I was thinking a MeOH+H2O mix might be more practical, but I have no data on (NH4)2SO4 or NH4NO3 solubility in those mixtures, so I have no way of knowing what would be a suitable mixture..

EDIT: I have finally been rewarded for wallowing in patents for hours.. I have attached several gems of solubility data.

Ethylene glycol is looking like a very good canidate. It's cheap and extremely accessible, and is not particularly flammable or volatile. I don't know for sure, but I would guess that (NH4)2SO4 has a similar solubility in EG compared to MeOH. Maybe 1g/100ml EG at most? Where NH4NO3 is going to have a solubility of over 50g/100ml EG when warm. That looks like it could very usable. The addition of some water would even further increase the NH4NO3 solubility, but I wonder if the solubility difference between it and (NH4)2SO4 would be better?

[Edited on 3-1-2009 by 497]

Attachment: solubility data.doc (134kB)
This file has been downloaded 554 times

not_important - 3-1-2009 at 22:48

Potential problem with water-alcohol (or polyol) mixes is a phase separation into water+salt rich and alcohol-rich layers.

It is possible that the solubility of ammonium sulfate in EG would be higher than in MeOH, should be easy to test.

Another difficulty with EG is what you list as a advantage - its low volatility. Once you've got NH4NO3 crystallising out of it, how do you dry them? The EG will hang around for some time, necessitating recrystallising from water at least once.

Is that pure EG, or an antifreeze which often has additives?

I would start with making a solution of the fertilizer, then adding enough CaCO3 or Ca(OH)2 to react with the sulfate plus a slight excess. Boil gently to drive off the released ammonia, adding water to keep the volume constant, until little ammonia is given off. Either move to a water bath or remove from heat and wrap the container in insulation to slow its cooling; let settle for awhile, then decant through a filter. Treat the cold filtrate with ammonium carbonate or NH3+CO2, heat to around 40 C, cool, filter. Bring filtrate to a boil to drive off excess of NH3+CO2, cool, decant-filter.

[Edited on 4-1-2009 by not_important]

NH4-NO3-SO4.png - 21kB

chloric1 - 4-1-2009 at 12:44

Quote:

Originally posted by not_important
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]

That sounds like a good system, especially if you can find the fishtank used at a garage sale or on craigs list. I only have 2 questions. What is a cartridge heater? And what volume of jar did you use to what volume of fish tank?

I have been wanting to further explore the diffusion techniqu for some time. Just have not had supplies or time to plan the setup since leaving it undisturbed at constant temperature for days or weeks is a little challenging but certainly possible.

My diffusion experiments would consist of precipitation by diffusion to form larger crystals. Like crystaline calcium hydroxide or barium thiosulfate. Just put small lipless cups of your saturated reagent solution with excess solid in the bottom and put both containers in a low form large container and carefully add water until the level covers both containers with several centimeters of water and pour melted paraffin over the water surface to block out dust and breezes and set aside for a few weeks. Ions should diffuse out of the beakers and deposit crystals of the least soluble component when solubility product is exceeded.

497 - 4-1-2009 at 14:19

Quote:

Once you've got NH4NO3 crystallising out of it, how do you dry them? The EG will hang around for some time, necessitating recrystallising from water at least once.

Couldn't you just wash it with something like petroleum ether or acetone to get the EG off?

Quote:

Is that pure EG, or an antifreeze which often has additives?

Some antifreezes are virtually all EG, a little bit of diethylene glycol, probably a small bit of water, and sometimes a corrossion inhibitor like sodium phosphate in small amounts, and of course the bright green dye (fluorescin?). Nothing that should interfere.

Quote:

I would start with making a solution of the fertilizer, then adding enough CaCO3 or Ca(OH)2 to react with the sulfate plus a slight excess. Boil gently to drive off the released ammonia, adding water to keep the volume constant, until little ammonia is given off.

I was really hoping to avoid dealing with CaSO4.... But if all else fails, I will go that route. Actually I have some Ca(NO3)2 fertilizer that will yield even more NH4NO3 when reacted with the sulfate... Maybe I'm overestimating the trouble this would be.. all I know is last time I had to deal with CaSO4 I vowed never to do it again.

EDIT: I tried it on a small scale today. I mixed 192g Antifreeze (straight out of the bottle) with 139g Fertilizer (straight out of the bag). I heated it in a boiling water bath for at least 30 minutes while stirring. It looked like a large portion of the solids dissolved. The surface of the glycol was covered with an ugly brown scum. Then I filtered it through a paper towel, which caught the scum nicely and left me a nice only slightly foggy solution. It worked pretty well, the filtering went quickly. The glycol solution gained about 57g (plus some that was lost as residue, etc). I took a small aliquot of the filtrate and left the rest in the water bath to slowly cool and hopefully give some decent crystals.

I set the aliquot (~2ml) in the freezer for a few minutes. It solidified with precipitated tiny needle crystals. I then let it warm back up to room temperature, where most of the precipitate redissolved. I then tried mixing in some toluene. As I not_important said, the two did not mix at all. So I poured off most of the toluene and let the rest evaporate. Then I tried MEK. Same thing happened. So I evaporated off the MEK and tried denatured ethanol (~10ml). This worked nicely, a substantial amount of fine white precipitate came out. Since AN is substantially soluble in ethanol using isopropanol is another possibility that would decrease the solubility even more. I'll have try it and see if they mix. If not, then a mixture of isopropanol and ethanol would probably be the ideal choice.

[Edited on 4-1-2009 by 497]

not_important - 5-1-2009 at 01:02

Quote:

What is a cartridge heater? And what volume of jar did you use to what volume of fish tank?

http://www.watlow.com/products/heaters/ht_cart.cfm

and roughly 1/2 to 1/3 the total volume of the tank. Smaller jar means more of the first load ends up in the solution, which is why you reuse it several times.

Quote:

Couldn't you just wash it with something like petroleum ether or acetone to get the EG off?

Not non-polar solvents like pet ether, basically if the solvent isn't moderately soluble in water then EG is likely only very slightly to insoluble in the solvent at room temperature or close to that. Acetone might work, but AN is slightly soluble in that so a tiny bit will be lost.

Quote:

I'd stick with the mixed ammonium nitrate/sulfate, as it has no other uses while Ca(NO3) can be used to make nitrates of metals with soluble sulfates; magnesium nitrate is useful for cleaning carbon and organic crud out of glassware and can be had by mixing Epsom salts and calcium nitrate.

A bit of boiling often coarsens the calcium sulfate. After that letting it settle for some time with occasional light tapping of the container, followed by decanting or siphoning of the supernatant liquid away from the precipitate, usually works well. Afterwards filter the liquid to remove small amounts of suspended sulfate, don't try to filter the bulk of the precipitate with conventional means.

To get the remaining liquor from the calcium sulfate, I dumped the gloppy precipitate into a tofu or cheese press lined with an outer layer cheese cloth and an inner (touching the precipitate) layer of muslin, calico, or cambric of plain cotton or linen. Have it set up over a plastic basin, scoop in the precipitate, drop on the top and weights, and go away for some hours or longer.

You can be frugal and wash the sulfate precipitate well, spread it in thin layers and allow it to dry, break it up into coarse chunks, and dry at 190-230 C to get a decent drying agent. If you keep at at 30C or cooler it's a good as H2SO4 for drying most gases, although it will absorb ammonia and small amines to some extent.

BTW - nice research on your approach. That data is around but buried in difficult to find places, so DIY is often the way to go.

[Edited on 5-1-2009 by not_important]

497 - 5-1-2009 at 15:16

Quote:

Yes I found that experimentally.. But I think washing with EtOH/iPrOH will do the trick.

The beaker of filtrate (from my above edit) has been sitting near room temperature for about 20 hours now. It has grown some amazing crystals on the bottom and lower sides. They consist of clear needles that are up to 1.5-2 cm long, and maybe 1 mm thick. Very pretty. Now I have it in the refrigerator (and later freezer) to see how much bigger they'll grow. I have to say I'm surprised, I expected to get a big gelatinous mass of fine crystals.

EDIT:

After sitting in the fridge for a few hours and then the freezer for a few more hours, this is what I ended up with:

The green tinge is from the dye in the EG residue.

[Edited on 5-1-2009 by 497]

chloric1 - 6-1-2009 at 08:28

not-important That is awesome! When I get back to the USA, I will call a watlow distributor. Nice touch with the mineral oil as aa heat dispersant. I might want to get some silicone oil from chemistrystore.com as I think it can work at higher temperatures if needed.

Just should mention that ammonium nitrate is appreciably soluble in 40% aqueous alcohol. Potassium sulfate is only 0.5% soluble in the same. That might be easier and cheaper than using straight methanol but I would like to try both.

497 - 6-1-2009 at 16:16

Woops, here's the picture I meant to post instead of the second one. Gives a little better view. Of course the edit function has expired...

KoiosPhoebus - 8-2-2023 at 17:29

Quote: Originally posted by DeAdFX

Quote:

Bit of an old thread, but I came across this while searching to see if anyone had written up my method for preparing ammonium nitrate from diammonium phosphate (https://www.sciencemadness.org/whisper/viewthread.php?tid=15...) and I thought I'd reply.

Yes, you can use the diammonium phosphate as dicalcium phosphate (calcium hydrogen orthophosphate) has very low solubility in water (0.018g/100mL). In fact it has lower solubility than calcium sulphate (0.26g/100mL).

I would not recommend using the monoammonium phosphate (aka ammonium dihydrogen orthophosphate) as it produces monocalcium phosphate which is somewhat soluble in water (1.8g/100mL). Some literature (e.g. https://pubs.acs.org/doi/abs/10.1021/acs.iecr.5b02100) even argues that monocalcium phosphate itself is much more soluble than that, it's just that some of it decomposes into phosphoric acid and dicalcium phosphate when dissolved.