ammonium nitrate--> nitric acid

[roux]

tho this may not be the most efficient or cost friendly way of obtaining nitric acid, i think its a method worth exploring. it uses ammonium nitrate and hydrocloric acid.

ammonium nitrate is made of a possitive ammonium ion(NH4+) and a negetive nitrate ion(NO3-)

hydrocloric acid is made up of a hydrogen ion just like the one in ammonium nitrate and a cloride(Cl-) ion that is more negetive than the nitrate ion.

i figure that if one were to mix the two compounds, the cloride ion would combine with the ammonium ion making ammonium cloride. this would result in a left over hydrogen and nitrate ion that would be forced to combine to make nitric acid.

HCl+NH4NO3--->NH4Cl+NHO3

i have not attempted this reaction, but i would be thankful if anyone were to tell me how it goes.

[Mendeleev]

Nice method, I have considered it, but it has a problem, unless you're using HCl gas your HCl will be at least 62% water, and it doesn't have the hygroscopic properties of sulfuric acid, so the acid you distill over will be low concentration, at most you'll get to the azeotropic point. Why do you say it is more costly? HCl= $2.50/gallon, H2SO4= $8.00/liter, NH4NO3=$6/25kg KNO3=$15/25kg. And now that I think of it, not so bad considering you could get azeotropic acid for cheap and then distill with H2SO4, which can be recycled. Eureka, that should save me a ton of money! One thing I am not sure HCl, NH4Cl, and HNO3 completely dissociate in solution (although NH4 doesn't NH4Cl does), so won't all the ions just be floating around in aqeous solution causing HCl to distill first since it has a lower boiling point than all the others?

[Marvin]

"i figure that if one were to mix the two compounds"

Nice try, but such speculation is almost always useless at your level of understanding (I cant think of a more tactful menthod of saying this). As you progress in chemistry you'll have a much better grasp of what can work and how.

Mendeleev is a little uncertain in his terms, but his basic idea is correct. What you get is a nasty mixture of ions and oxidised compounds, Aqua Regia, and a method for obtaining nitric acid would have to involve extracting just this from it. It would be uneconomical for any method suggested before, or that I can think of. I'm sure this topic has come up before though perhaps not in this section.

[vulture]

NH4Cl <-----> NH3 + HCl is an equilibrium and will greatly interfere.

[gliper]

If it leaves NO3 or HNO3 in any resonable concentration you should be able to make uria nitate (mix with uria) and from there nito-uria (7000 m/s). Both more powerful and sensitiv than AN.

As for HCl being cheaper when you can get ether one its not that simple. $2.50 is for 30% where as sulferic Is 30% for $3 a gallon (at auto shops I've read) or 90ish% $17 a gallon, also remember 1 SO4 is worth 2 HCl s. I calculated that the cost of HNO3 is the same for my drain cleaner and HCl but SO4 is more consentrated and easyer to use.

[THEBOARDER]

ok.
i have fertilizer 10-0-30 , 30% k2o and 10% N , name ( N-NH4 ) , can i have nitric acid from this

[Rhadon]

THEBOARDER, you're having three posts, and every single of them is nothing but crap. HED.

[tom haggen]

Since this is the official NH4NO3--> HNO3 thread, I have a couple quick questions relating to this topic. For starters, how much more efficient is KNO3 and H2SO4 in the manufacture of HNO3, versus using NH4NO3 and H2SO4. Also If you did NH4NO3 and H2SO4 would the ratios be about the same as they are when using KNO3 and H2SO4.

[Boomer]

My compi crashed during 'reply', and my reply did no come up after logging in again. Sorry if it comes twice for some reason later:

Why should KNO3 be better? AN contains 20% more nitric and is much cheaper! And for the ratios: SEARCH or (even better) calculate. 1 mol H2SO4 converts 2 mol AN to 2 mol HNO3, leaving 1 mol (NH4)2SO4 over. H2SO4 is 98g, AN is 80g, KNO3 is 101g per mol. Now it's your turn....
P.S. If you use more (max. twice as much) SA, the process is cooler/safer and the NA decomposes less, making it more concentrated (less water and less NOx/nitrous acid).

[markgollum]

About reacting hydrochloric acid with ammonium nitrate.
In my experience, I have found that when ammonium nitrate is dissolved in 31% hydrochloric acid, crystals of what should be ammonium chloride precipitate (I say "should" because the crystals were not tested).
The precipitate quickly dissolved when more water was added, being completely dissolved when approx 75%IIRC of the original acids volume of water was added.

With this knowledge in mind, this is how I think nitric acid could be produced from HCl(aq) and AN (with terrible yields).

(1) Dissolve AN in conc hydrochloric acid.
(2) Filter out ammonium chloride.
(3) Boil soln until either more chloride ppts, (unlikely because of increased solubility of chloride due to high temp, and NH4CL(aq) > NH3(g) + HCL(g))
or you reach azotrope, (if ammonium chloride ppts then refilter.
(4) Cool solution to try to ppt more chloride. (saves on H2SO4 by reducing this from happening H2SO4 + NH4Cl > HCl + NH4HSO4)
(5) Extract nitric acid from solution with either sulfuric acid and distillation, or by using a dichloromethane extraction.

I am sure that the yields would be abysmal, but considering the value of nitric acid, and the availability and cost of reactants, it might just be worth it.

[simply RED]

How do you thinik (NH4)2SO4 can be reached?
The reaction stops when you have (NH4)HSO4 or KHSO4 it does not matter.

[Marvin]

You can almost certainly make an aq mixture of HCl and HNO3 where the HNO3 is the larger componant. But if remotly concentrated, it will be highly reactive (destructivly oxidising, etc) and removing the HCl will be either tricky or expensive. Standard method is to distill from silver nitrate, producing also silver chloride which is difficult to recycle, and needs surprise surprise, nitric acid free from chloride. Overall it will certainly make dilute nitric acid, but its a hell of a lot of work and you then have to concentrate it.

Steps 3 and 5 in the suggested process look unworkable to me and you cant distill nitric from nitric+AgNO3+H2SO4 in order to concentrate and remove chloride at the same time - just in case that is what you were thinking of suggesting next :)

[tom haggen]

I would just like to follow up on my previous post. I did further reasearch on this topic and found that it is extremely dangerous and unefficient to manufacture HNO3 with NH4NO3 and H2SO4. As for manufacturing HNO3 with hydrochloric acid and ammonium nitrate, it might be safer I don't really know. It seems pointless unless you want to manufacture some HNO3 with low concentrations.

[grendel23]

Tom,
On what do you base your statement that using NH₄NO₃ to make HNO₃ is "extremely dangerous and unefficient".
I have read posts from members who had poor yields, but I had not heard that the procedure was more dangerous than using other nitrates.
In this thread, http://www.roguesci.org/theforum/showthread.php?t=2333, Rosco Bodine states;

"For making azeotropic nitric acid 68 per cent which
distills at about 120 Centigrade, Ammonium Nitrate
is actually preferred . There is added to the mixture
sufficient water to account for the water to be contained
in the azeotropic nitric which will result , plus a few
more per cent of theory because the dehydrating action
of the sulfuric tends to disrupt the azeotrope in favor
of more concentrated nitric being produced . A nearly
100 per cent yield of azeotropic nitric acid 68 per cent
is the result from a very smooth and rapid distillation
where the residual mixture remains a liquid
having no precipitated acid sulfate salts at the
120 degree C temperature of distillation."

I have not tried this yet myself, but I have a lot of respect for the knowledge and experience that Rosco has shown.
I am going to give this a try soon, I like the idea of no fudge in the flask afterwards.

[tom haggen]

"Ammonium nitrate, unlike alkali metal nitrates, breaks down at a relatively low temperature. The reaction may be self- perpetuating once started."

I'm quoting polverone from a different forum. However, it was uncertain if this was 100% accurate. Also it was stated later on in the same thread that your final product would be heavily contaminated with ammonium nitrate.

Another quote from the same thread was "Heating this a little too hot is a serious problem since over a cetain T HNO3 will evolve NOx that will destroy your NH4NO3 quite fast (maybe explosively) and your HNO3 also!" by PHILOU Zrealone

[markgollum]

Marvin suggested that nitric acid with HCl in it would not be useful to preform nitrations with.

You can almost certainly make an aq mixture of HCl and HNO3 where the HNO3 is the larger componant. But if remotly concentrated, it will be highly reactive (destructivly oxidising, etc) and removing the HCl will be either tricky or expensive

While I was thinking about this, I noticed that the only reason that the this equilibrium is on the right, is because the chloride is being removed from the system.
HCl(aq) + NH4NO3(aq) > HNO3(aq) + NH4Cl(s)
HOWEVER, (this is the interesting part) once the amount of chloride that precipitated is removed, the equilibrium should shift again, (way to the left) because of the excess nitric acid, and the fact that nitric acid is a stronger acid than HCl(aq).

excess HNO3(aq) + remaining NH4Cl(aq) > HCl(aq) + NH4NO3(aq)
(pretty much completely on the right)
Then in step 3, (boiling) the hydrogen chloride would evaporate pushing the equilibrium even further towards NH4NO3.
Because of this, I think that the remaining solution would contain only AN, nitric acid, and water with only an insignificant amount of HCl(aq).

[Guerilla]

Take into account that raising the temperature will change the whole equilibrium constant, and most likely so that it will go more towards HCl + NH4NO3 due to it being an endothermic direction. Let alone how it will also favour the decomposition of NH4Cl and oxidation of HCl as has been stated earlier. While this may not be an issue in a small scale its good to note that Cl2 that forms (aqua regia) can react with NH4Cl forming nasty NCl3.. This method seems to be of no practical use for making HNO3, mostly because of the equilibrium reactions that interfere with each other.

[Boomer]

Maybe you are right. But I remember reading somewhere that if heated stronger, the sulphuric is completely used up till the X2SO4 is formed.
You should avoid this by using a little over 1 mol acid (not 1/2 mole) per mol of nitrate as the NA decomposes faster at the temp needed for full conversion.

I am no chemist but I could imagine that, even though the hydrogen sulphate is the weaker acid, it converts more AN to NA. Perhaps because NA is more volatile? Even though HCl is stronger than SA, it is driven of if SA is heated with NaCl.

To tom haggen: Using NH4NO3 IS efficient! 1 mol AN give 41ml HNO3 in theory, and that is what you get. I usually stop at 37ml as then the dropping gets slower at the same heat, and I want to avoid decomposition.
The resulting NA is nearly white (only very faint yellow), and works well even for RDX. :)
And it is NOT that dangerous. Pure AN can decompose explosively if heated above 200C. But it is mixed with 2.5 times it's weight of SA, and most is converted to the non-explosive sulphate anyway. The rest cannot heat the whole mass enough, and you work slowly at low temp anyway to avoid diluting your nitric with water from its own decomposition.
I have done it lots of times, and it stops boiling if you turn the heat of, even in litre batches. :D

[tom haggen]

I'm not a chemist either, and I'm somewhat skeptical of how dangerous this process is. However, something I don't like to do is waste my time on projects that aren't going to give me any results. This method of manufacturing HNO3 is completely unefficient. As madscientist said "It technically will work, but not well at all. The ammonium hydrogen sulfate formed will be decomposed into sulfuric acid and ammonia from the heat; the ammonia gas will then react with your nitric acid to form ammonium nitrate. You will get some nitric acid, but I doubt that your yields will be worthwhile, for much of the nitric acid will have been converted back into ammonium nitrate, as well as the fact that what nitric acid you do have will be heavily contaminated with ammonium nitrate." With this happening I find it hard to believe that you manufactured HNO3 capable of synthesizing RDX.

[Rosco Bodine]

There is a German patent DE280967 which should provide
some accurate information about making nitric acid from
ammonium nitrate and sulfuric acid . Maybe Rhadon
or someone else here can translate the German to English .

I know that what I reported concerning azeotropic nitric
is accurate .

And yes ammonium nitrate does decompose at lower
temperatures than metal nitrates . However , unless
that decomposition was highly catalyzed by certain
impurities , ( chloride comes to mind ) , nitric acid
is formed and distills at a fairly low temperature ,
which is well below the temperature at which there
would usually occur significant decomposition of
ammonium nitrate .

In COPAE pg 370 , Davis mentioned that ammonium nitrate in
a large excess of concentrated sulfuric acid produces a mixture
which contains free nitroamide , but that the nitroamide
is hydrolyzed to nitric acid if the mixture is "digested"
for a time with gradual and gentle heating at 90 to 120 C
which is precisely the way a nitric acid distillation proceeds .

"Failed experiments" are not a waste of time , for you
are then equipped with good information about what doesn't
work in reality as it would seem it should work in theory .
Experimental results beat the hell out of anyone's guesses
and speculations based upon theory .

[Rhadon]

Rosco, I was unable to find this patent on the net. Can you supply me with a link or otherwise give me access to the file? Then I could take a look at it.

[Rosco Bodine]

Attached is the zipped pdf for DE280967 .
For some reason the file size upload limit
was exceeded for the uncompressed pdf file .

Go to the following page and enter the de280967 in the
second search box , then hit the Go button and follow the links .

http://ie.espacenet.com/espacenet/ie/en/e_net.htm

or a couple of direct links which may or may not work ,

http://l2.espacenet.com/espacenet/bnsviewer?CY=ie&LG=en&DB=EPD&PN=DE280967&ID=DE++++280967A++I+

or try clicking the highlighted patent number on this page

http://l2.espacenet.com/espacenet/viewer?PN=DE280967&CY=ie&LG=en&DB=EPD

[Rhadon]

Thanks for the patent, Rosco. I've been to the German version of the page you linked me to, but for some reason I wasn't able to get this patent and I was told that it wasn't available on the server. Probably I did something wrong.

Anyway, here's the translation. I begin with page 1, line 25 as we don't need the preamble. Just remember that the patent is from the beginning of the 20th century.

There's no doubt that ammonium nitrate can be reacted with sulfuric acid just like sodium nitrate. But applying the same process would result in considerable losses in the nitric acid yield because of the height of the temperature required. According to Berthelot, ammonium nitrate does already melt at 152 °C, but decomposition starts at the same temperature if the reaction is carried out in an iron vessel. At 185 °C the salt will decompose completely under the evolution of N2O.

Now the invention is to measure the amount of sulfuric acid so that the production of nitric acid is already completed below the temperature reaches 152 °C or 185 °C respectively. This can easily be achieved by using so much sulfuric acid that for each molecule of ammonium nitrate there is at least one molecule of sulfuric acid. Hereby the reaction product does already melt completely and the resulting nitric acid has already completely distilled over without any ammonium nitrate being decomposed, at temperatures that don't exceed 120 °C significantly.

Important advantages of this process are:
A higher, almost theoretic yield of nitric acid.
Significantly lower content of nitrous acid in the nitric acid produced.
The distillate is free from sulfuric acid.
Calm way of working, since the reaction mass doesn't show the tendency to rise in the reaction vessel during any time of the process.
The possibility to use this process for concentrating less concentrated and impure nitric acid, such as that gained with the "air-nitrogen process" [probably reacting NH3 with air at elevated temperatures].
The possibility to neutralize the remaining excess of sulfuric acid with ammonia and thereby gaining commercial grade ammonium sulfate what cannot be done to the same extent with sodium sulfate when sodium nitrate is used respectively.

[Rosco Bodine]

Rhadon ,

Thank you for the translation .

That's good information . Generally I have found that reports
from German chemists are highly reliable .

Prills and Drain Cleaner , yeah ....it doesn't get any easier :)

Just bring up the heat slowly and use a slight excess of sulfuric
for satisfying the equation

H2SO4 + NH4NO3 -----> NH4HSO4 + HNO3

The water in the 92.5 per cent sulfuric acid drain cleaner
actually helps the reaction in that it helps hydrate any nitroamide
formed to convert it to nitric acid . Also , the highly soluble
ammonium acid sulfate by product remains a liquid , not
complicating the distillation with any bubbling curd of solids ,
and probably also contributes an affinity for water to keep
it from distilling over at lower temperatures . Ammonium salts
also have the property of reacting with nitric oxides to decompose
them to nitrogen and water , similarly as does urea .

[Boomer]

Well, obviously Madscientist was wrong and those of us that get 90% yields of RDX-grade nitric acid from AN (like me) were right. The patent gives a scientific explanation why it works! :)

In another book I found the statement that at SA : NA : H2O = 1:1:1 (mol), the formation of N2O5 is just prevented. With less water it starts to form, decomposing to N2O4/NO2 and O2. This ideal ratio is 55% SA, 35% NA and 10% H2O (weight). This means drain cleaner (92%) is more than strong enough, even if the AN is a little moist.

Another thing I found: If you use a little more SA, you first destil RDX-grade NA off, then add more AN to fully use up the SA (if it is your most expensive precursor). The second crop is less concentrated and contains more NOx (deeper yellow instead of nearly colourless), but can still be used for MNT/DNT etc. where NOx and a little water do not matter much.

[NTAN]

Some thing here disturb me,
I know that HCl and strong oxidizers react and emit Cl2

so why you got HN4Cl and not free Cl2?

[Boomer]

We don't get Cl2 because there is no Chlorine in the reaction! The problem is rather that some of the nitric acid is decomposed and forms dihydrogen monoxide ..... :p

OK, you might have missed that the conversation moved towards SULPHURIC acid plus ammonium nitrate.

P.S. What is HN4Cl ?

[0EZ0]

He most likely meant NH4Cl or Ammonium Chloride.

I don't understand all this who-ha about the credibility of producing Nitric Acid from H2SO4 and NH4NO3. It should be well noted that most Ammonium salts tend to decompose at relatively low (go figure) temperatures as compared to other more traditional ionic salts. An easy way to avoid substantial decomposition and thus contamination of your final acid is to as noted distill at low temperatures with an excess of H2SO4. It makes quite alot of sense does it not?

I suspect what madscientist was getting at was the trend of low temperature decomposition of Ammonium salts. Distilling at such high temperatures as 120 degrees Celsius will also result in alot of NOx contamination of the final acid and so is a less suitable temperature for distilling. People who do distill at higher temps are probably more impatient and think that a little contamination is worth the shortened distilling time. There is nothing wrong with doing that with methods using KNO3, NaNO3 and Ca(NO3)2. You really only have to be wary of the Sulfate salt produced and it's decomposition when using methods incorporating NH4NO3.

Note that most Ammonium salts (and alot of other compounds) don't just "go poof" at their written temperature of decomposition. Decomposition starts at much lower temperatures and is increased as the temperature rises. This should already be known but I thought for the sake of clarity I would add it.

[nbk2000]

Fortunately for us all, madfag and his delusions of "kHemIcL" knowledge, have long since been purged. :)

Chemistry is not simple 1+1=2, try as you might, because there's many variables involved that are not linear progressions from point A to point B. He couldn't grasp that concept. :rolleyes:

Thus, people who've empirically determined that the process works, by actually trying it, have knowledge on the subject that cook-book theorists will never have. :p

[tmp]

I've always made NA with KNO3 and SA. I can see how AN would work with
SA, but even the most concentrated HCl(liquid) is still mostly water. Even
though it's cheaper, HCl looks like too much work for too little acid. I'll try SA
with AN method to satisfy my own curiosity.

[ikbendirk]

How about the NH4NO3 -> HNO3 + NH3 rxn?
It is a self-sustaining reaction I believe*, but there are only two problems:

- The reaction should not be mistaken with the rxn which produces nitrous and water, I believe the HNO3 rxn can be catalysed by Cl- ions.. am I correct?

- As a product, you will most likely get two gases, namely NH3 and HNO3 as a gas (because of the heat) one must try to condensate the HNO3 gas as quick as possible to get it as liquid. The concentration should be very high, as no water is included with the reaction...

[HNO3 = Boiling point: 121 C (69% boils at ca. 86C)]
[NH3 = Boiling point: -33.3 C]

What do you say.. could this be practical?

[Edit:]
* according to "Basic Inorganic Chemistry" the ΔH of the HNO3 rxn is 171 kJ/mol, meaning the rxn is endothermic and cannot possibly be self sustaining. I quote:

Many ammonium salts volatilize with dissociation around 300˚, for example:

NH4Cl (s) = NH3 (g) + HCl (g) ΔH = 177 kJ/mol
NH4NO3 (s) = NH3 (g) + HNO3 (g) ΔH = 171 kJ/mol

Salts that contain oxidizing anions may decompose when heated, with oxidation of the ammonia to N2O or N2 or both, for example:

(NH4)2Cr2O7 (s) = N2 (g) + 4 H2O (g) + Cr2O3 (s) ΔH = -315 kJ/mol
NH4NO3 (s) = N2O (g) + 2 H2O (g) ΔH = -23 kJ/mol

<...>

Nitrous oxide, N2O, is obtained by thermal decomposion of molten ammonium nitrate:

NH4NO3 (s) --250˚--> N2O (g) + 2 H2O (g)

I don't know quite how to seperate the N2O and HNO3 reactions, but i guess it should have something to do with the temperature which should be kept above 300˚.

I am still a bit puzzeled about the '=' in the equitation, but I guess it is meant as an '-->'.

Well I guess it would work as follows:
1) Heat the NH4NO3 (300˚+)
2) Collect the gases
3) Cool the gases down to any temperature between 80˚ and -30˚ this could be done with ice, or just plain water.
4) let the NH3 gas escape and you should have some nice pure nitric acid.

http://www.redrival.com/blackstorm/HNO3.txt

I have not tested it yet, but I guess it should work...

[nbk2000]

Heating ammonium nitrate gets you nitrous oxide (laughing gas) and water, not ammonia and nitric acid.

Chemistry is not 1+1=2.

Just because AN is made from ammonia and nitric acid doesn't mean you'll get the same two chemicals out of it when you break it down. :rolleyes:

[Boomer]

If you want to learn about NA decomposition, read the Escales book on AN explosives. He gives 7 (IIRC) ways of reaction, one of them being into HNO3 + NH3, others give N2O + H2O; NO + N2 + H2O; N2 + H2O + O2; NO + NO2 + H2, ... I forgot some.

You can get every possible gas mix, but in the end he says that it comes down to either laughing gas and water, like NBK said, or nitrogen, oxygen and water (if fully detonated).

I would not try to get high on the N2O though, because you might have traces of NH3, HNO3, HNO2, NO and NO2 in it ... :p

[K9]

Although there are going to be a variety of gases produced during the decomposition of NH4NO3 (as stated by the above poster), the main gases formed are N2O and H2O thus making that quite an ineffective method of production.

As far the N2O goes, it's not too hard to purify it by bubbling through cold water which will remove the nastier of the products. Then again, I wouldn't try inhaling it.

[nbk2000]

I would not try to get high on the N2O though, because you might have traces of NH3, HNO3, HNO2, NO and NO2 in it ...


TRACES being the operative word in the above.

You might get nitric from AN, but it wouldn't be under conditions you could duplicate with standard glassware, likely requiring high pressures and temperatures, same as those required to create it from air in the first place.

And there's 9 forms of ice crystallization possible, but only 1 happens in your fridge, so what's your point?

[K9]

Not really on topic, but I'm fairly certain that there are even more than 9 forms of ice. Something like more than 12 even. Which again reinforces the point you're not going to have an effective production method this way.

[TOX]

Hell ya I got the video at last ammonium nitrate --> nitric acid .English Subtitles for at least the concentrations and temperature would be great.

http://rapidshare.de/files/28607245/_________________.wmv.html

Enjoy,guys

[nbk2000]

I love how the raghead chemist is wearing latex gloves and wearing sandals while handling open beakers of Sulphuric acid and fuming NA. :rolleyes:

The still they used in the video would be easy to duplicate. An improvement would be to put a U bend in the acid outlet tube to keep uncondensed acid fumes from venting out. Not only is that wasteful, but also hazardous to the operator.

As for ratios, concentrations, and yields, UTFSE here and you'll find plenty of discussion on these details in other threads. :)

[NoltaiR]

I also really like that they use some proper glassware.. but they also don't forget the necessary cooking pot :-P oh there's nothing like watching the middle-easterners play with chemicals..