Sciencemadness Discussion Board

nitric acid from calcium nitrate

497 - 27-1-2008 at 13:21

I recently acquired 50 pounds of calcium nitrate fertilizer. Later to my annoyance i found out it was actually a calcium nitrate ammonium nitrate double salt. But that's the a big problem, theres only ~7.5% ammonium nitrate, I just dissolve it along with CaOH and boil off the NH3.

What I wanted was a nice easy nondistillation way of getting HNO3, which calcium nitrate seemed perfectly suited for, considering the very low solubility of calcium sulfate. It was the only nitrate i could get besides KNO3 (which i did get 50 pounds of too.)

The real problem is the consistency of the damn calcium sulfate. After eliminating the ammonium nitrate I boiled the solution down to a very thick sludge and mixed it with ~65% sulfuric acid (boiled down from the gallon of battery acid I just aquired.) It forms a nice thick paste of gypsum-nitric acid mix the seems to defy most metods of separating. Of course I could distill it but thats what i was trying to avoid in the first place. So my question is can I separate the mixture and if so, how? Short of distillation or dilution...

The feed store i got the fertilizers from does carry sodium nitrate in the summer but thats a couple months away. The stuff I managed to get was leftovers. Also I would love to hear anyone's suggestions of other interesting things to do with nitrate fertilizers.

[Edited on 27-1-2008 by 497]

microcosmicus - 27-1-2008 at 14:12

One thought is to figure out a way of slowing down the process so as
to produce bigger crystals which could easily be filtered out. That, of
course, could be done by dilution, but that method is ruled out. Maybe
try running the reaction at a low temperature to slow it down and add
the acid as slowly as possible. The sludge of Ca(NO3)2 ought to
have a low freezing point and 65% sulfuric acid freezes at something like
-40C. Also, at low temperature, the stuff should be quite thick, further
reducing the rate of mixing.

Another fact worth noting is that the solubility of CaSO4 increases when you
chill the solution. Maybe you could cool down your mess to dissolve some of
the fine precipitate and gradually rewarm it to form bigger crystals.

Magpie - 27-1-2008 at 14:30

I have successfully made phosphoric acid from Ca3(PO4)2 and sulfuric acid by the principle you are attempting to use. I kept the reaction mix dilute until I could filter off the CaSO4 using a Buchner suction filter and Celite (diatomaceous earth) as a filter-aid. Then I evaporated the dilute (and clear) phosphoric acid to the strength that I wanted. This procedure should work well for making azeotropic (65%) nitric acid also.

[Edited on by Magpie]

not_important - 27-1-2008 at 20:41

I done the "fairly dilute solutions, filter" route, but always finished with distillation to get the constant boiling acid.

One reason for this is the solubility of CaSO4, which is enhanced by the mineral acid:

CaSO4 solubilities.gif - 34kB

497 - 28-1-2008 at 20:41

I think I solved the problem. Actually bypassed it, by simply dissolving the Ca(NO3)2 in water and adding the appropriate amount of Na2CO3 to yield NaNO3 which can then be used to make the nitric and the sulfate will be much easier to separate. I can't believe i didn't think of this earlier. I've got a pretty big batch cooking, it will be so nice to finally have a decent amount of nitric acid :) $25 for a bag of fertilizer that should yield something like 14 kg of HNO3. I'm happy.

[Edited on 30-1-2008 by 497]

497 - 30-1-2008 at 17:41

So something a little interesting happened today. I mixed 1 mol Ca(NO3)2 , which has 1/5 mol NH4NO3 in it with 1 mol Na2CO3 and added 2 liters of RO water. It turned very white with CaCO3 particles as i expected, but it also bubbled off quite a bit of CO2, which surprised me. The bubbling lasted quite a while, not all at once. The reaction i thought was going on was:

Ca(NO3)2 + Na2CO3 --> CaCO3 + 2NaNO3

i don't see why CO2 would be evolved. I was thinking the NH4NO3 would not be effected, and there was no ammonia smell. Maybe some other impurities in the fertilizer? I'll mass the leftover CaCO3, maybe that will tell me something. I'm probably missing something obvious...

chemoleo - 30-1-2008 at 18:08

I suppose these fertilisers tend to be slightly acidic, so the formed ammoniumcarbonate can decompose into another NH3 salt and CO2. Same for the Na2CO3.
Best is, check the pH of the Ca(NO3)2 solution and see if it is below 7, if so, then you have your answer.

aliced25 - 25-5-2013 at 00:46

Mate, if that is fertilizer, you are going to have to purify it before you do anything. Dissolve, filter, crystallize, dissolve filter, recrystallize - a giant pain in the ass, but that is why people buy pure reagents when they can. I recall having to lose approx. 1/2 of a bag of ammonium hydrogen phosphate during the purification process. It still worked out cheaper than ordering from a chemical supplier, but it was a weeks work.

blogfast25 - 25-5-2013 at 04:02

Here's a really thorough video on CAN and making nitric acid from it but it does involve distillation:

http://youtu.be/7akk5ppJjEw

Fantasma4500 - 25-5-2013 at 14:45

indeed i do..

Ca(NO3)2 + (NH4)2SO4 > 2NH4NO3 + CaSO4

meaning per mole of Ca(NO3)2 used you SHOULD end up with 160g AN relatively pure

ammonium sulfate is easy to get..

and if you mix up the NH4NO3 with H2SO4 you get........ yes you get (NH4)2SO4
so you can repeat it over and over and over (:

and you can use the AN ofcourse

you could also try vacuum filtration of it.. i suggest making a flat piece of nitrocellulose in which the HNO3 wouldnt react with, and would be able to filter off the CaSO4.. perhaps just take it all into a piece of cloth with some gloves on squeeze it into a container and decant it if needed..?

Beginner - 20-3-2016 at 05:29

Quote: Originally posted by 497  
I think I solved the problem. Actually bypassed it, by simply dissolving the Ca(NO3)2 in water and adding the appropriate amount of Na2CO3 to yield NaNO3 which can then be used to make the nitric and the sulfate will be much easier to separate. I can't believe i didn't think of this earlier. I've got a pretty big batch cooking, it will be so nice to finally have a decent amount of nitric acid :) $25 for a bag of fertilizer that should yield something like 14 kg of HNO3. I'm happy.

[Edited on 30-1-2008 by 497]

Sorry for a probably stupid question, but how does one do to remove all the Na2(SO4) from the nitric acid? CaSO4 has a very little solubility in water but Na2(SO4) dissolves with 47,6 g/l (0 C). Do you distill the nitric acid anyway? I prefer the nondistill-method because it's much moore safe, but don't understand how to remove all the sodium sulphate.

[Edited on 20-3-2016 by Beginner]

Metacelsus - 20-3-2016 at 08:05

Yes, distillation is required. If properly carried out, distillation of nitric acid isn't very dangerous, especially if you can use aspirator vacuum and keep the temperature low. Just make sure you have good ventilation.

Beginner - 21-3-2016 at 02:50

Quote: Originally posted by Metacelsus  
Yes, distillation is required. If properly carried out, distillation of nitric acid isn't very dangerous, especially if you can use aspirator vacuum and keep the temperature low. Just make sure you have good ventilation.

Ok, thanks! But it is a pity that it was so hard to remove the calcium sulfate with the first method , it would have been so much easier to do so.

ave369 - 21-3-2016 at 04:49

What's so scary about distilling nitric acid? If you have an all-glass apparatus or a retort, it can be done easily. And if you do not overheat the pot but heat it gently, no NO2 will emerge from the apparatus.

AJKOER - 21-3-2016 at 09:03

On a small scale one could produce high purity nitrate salt, or HNO3 without distilling if one has access to scrap copper.

For example, apply H2SO4 to your nitrate source but not in excess. Apply Cu to the mix and depending on concentration, expect NO or NO2 gas. For a nitrate, apply NO2 to the metal oxide (source: see https://en.m.wikipedia.org/wiki/Nitrogen_dioxide). Interestingly, applying NO2 to a metal iodide can result in the nitrite (same reference).

For safety, the NO2 generating vessel could be encased in a sealed larger vessel containing the metal oxide. Near the end of the process, one could introduce moist Na2CO3 to remove residual NO2. Nevertheless, recovering the created nitrate could still entail some NO2 exposure.

For HNO3, collect NO2 (or NO) in a vessel containing water and O2 (from, say, dilute NaOCl + H2O2). I would expect that exposure to sunlight would be catalytic as the intial reaction with water forms HNO3 and HNO2. As nitrite photolysis can create hydroxyl radicals (reference: see eq. [2] and [4] at
https://www.google.com/url?q=http://scholar.google.com/schol... ), the following radical reactions:

NO + .OH + M = HNO2 + M
NO2 + .OH + M = HNO3 + M (Source: see reaction [13] at http://www.fraqmd.org/ozonechemistry.htm )

could promote gas conversion to acid.

The embodiment is likely on a small scale only due to safety concerns (limiting exposure to corrosive/toxic NO2). [Edit] The goal should be to present less risk than one would experience with the quality of ones distillation equipment.

[Edited on 21-3-2016 by AJKOER]

[Edited on 21-3-2016 by AJKOER]

hissingnoise - 21-3-2016 at 09:15

Quote:
The goal of the particular embodiment would be to present less risk than one would experience with a distillation path, if possible.

WTF! Risk? What risk, where?

AJKOER - 21-3-2016 at 09:34

Hissingnoise:

The level of risk likely increases with the quality of the equipment, as we are working with home chemists, from whom I recall pictures of some rather make-shift setups for distillation of wood to methanol, as an example. When trying to run a synthesis on a relatively larger scale (part or whole bag of fertilizer, for example), something other standard lab equipment may be more of an option.

Also, Nitric acid does attack rubber, tape,.. and may contribute to a leak.

But, I agree, performed correctly in a lab, limited risk, and I have edited my comment.

[Edited on 21-3-2016 by AJKOER]

hissingnoise - 21-3-2016 at 13:47

Quote:
I recall pictures of some rather make-shift setups for distillation of wood to methanol, as an example.

Rather a poor analogy I think, AJ ─ if I were to pyrolyse wood, my setup would look pretty improvised too, as any working temp. above 200° rules out my doing it in glass!

I mean, NA comes over @120°, ffs . . .