garage chemist - 12-11-2004 at 09:21
I had an idea about a new way to make nitric acid which doesn't involve sulfuric acid.
When NaHSO4 is molten with NaCl, gaseous hydrogen chloride is evolved. I've used this reaction to make HCl gas before (for the HCl fountain, a
classic lecture demo, works like the ammonia fountain).
Now, what do you think what happens when NaHSO4 is molten with NaNO3 or KNO3? Actually, fuming nitric acid distills from this mixture in good yield!
NaHSO4 + KNO3 ---in molten state---> HNO3 + NaKSO4
No need for H2SO4!
Axehandle, this might be of interest for you.
NaHSO4 can be easily bought OTC as "pH down" from the pool supply store.
Theres only one problem in this preparation: due to the high temperatures in the molten salt mixture, a considerable portion of the HNO3 decomposes,
and some NO2 is produced and the resulting HNO3 is most likely not over 80%.
You also need a ground glass distillation apparatus.
Was I the first to come up with this idea or has this method been described before?#
I haven't tested it yet, but I have distilled HNO3 from a mixture of 1mol H2SO4 and 2mol NaNO3 before. After the HNO3 has distilled off, the
mixture melts and gives a second yield of red HNO3 with lower concentration (~80%) and more NO2.
[Edited on 12-11-2004 by garage chemist]
axehandle - 12-11-2004 at 09:36
I suppose the NaHSO<SUB>4</SUB>-KNO<SUB>3</SUB> mix wouldn't, by a pure stroke of luck, be eutectic...?
Heating KNO<SUB>3</SUB> to melting point would make me nervous, since it melts at 334C and decomposes at 400C. Nitrogen oxides filling the
room would be nasty...
garage chemist - 12-11-2004 at 09:59
Only the NaHSO4 has to melt, the KNO3 doesn't need to be molten.
Molten NaHSO4 acts like conc. H2SO4 in a lot of reactions.
The NaHSO4- NaCl mix also produces HCl as soon as the NaHSO4 is molten. If the NaCl had to be molten, you'd have to heat it to 800°C, which
isn't necessary.
vulture - 12-11-2004 at 10:57
I think we've had quite a developed thread about this subject already.
garage chemist - 12-11-2004 at 12:28
I searched for "NaHSO4 HNO3" and it came up with zero hits.
But I searched again (after I created this thread) and it still comes up with zero hits!
There must be some kind of software bug, since it doesn't find this thread.
If a thread about this subject already existed, it must be really old because I've read most of the threads here for a long time (I've been
reading this forum for about a year before I registered)
Theoretic - 12-11-2004 at 12:58
Axehandle, KNO3 doesn't decompose to give nitrogen oxides at 400 C, it gives nitrite and oxygen instead. Only at 800 C - 1000 C does it decompose
to K2O and nitrogen and oxygen (or nitrogen monoxide/a bit of dioxide and oxygen), and even then only slowly.
Decomposing NH4NO3 catalytically with chloride inos as the catalyst would give you water vapour, nitrogen gas and nitric acid. I think this might be a
good method, however your HNO3 will be dilute.
Eclectic - 14-11-2004 at 12:36
Actually, NH4NO3 with chloride thermally decomposes to H2O and N2O (Nitrous Oxide).
blip - 14-11-2004 at 17:31
I have done this a few times, but not actually collecting enough nitric acid to do more than fume and dissolve the glue from a cotton swab that holds
the cotton on the end. On the other hand, seemingly half the time I'm in my lab I do the HCl producing reaction. In fact the idea for these
acid distillations came from <a href="http://mattson.creighton.edu/Microscale_Gas_Chemistry.html" target="_blank">a gas
microchem site</a> where they make HCl in the same manner. One time I had calcium acetate on hand from chalk + vinegar so I tried distilling
it. It worked quite well and yielded a few drops inside the test tube using my method. My method employs a medical clamping instrument that
resembles scissors to hold the test tube by having me hold on the "blade" end and the test tube being held where the metal steps are
adjacent to the rings for your fingers. The metal acts as a heat sink so sometimes condensations can occur above the holding point on the tube.
Anyway, I was bold enough to put up the swabbed sample up to my nose and actually SMELL it without wafting. Heh, that really burned and I jerked
backward in my seat. Btw I use sodium bisulfate instead of sulfuric acid because I have yet to locate a source for the latter.
Theoretic - 15-11-2004 at 10:39
AN decomposes to N2O without catalyst or with most catalysts. With chloride it decomposes to nitric acid, nitrogen and water.
Eclectic - 15-11-2004 at 15:07
I don't have a reference handy, but in a search for viable methods of producing N2O, I often read that Cl- was a catalyst for the production of
clean N2O with LESS contamination by other N oxides. I haven't actually tried it though. I'm a bit afraid of runaway decomposition
(explosion)
Theoretic - 16-11-2004 at 11:39
Here.
Eclectic - 17-11-2004 at 06:17
http://mattson.creighton.edu/N2O/
Theoretic - 17-11-2004 at 07:35
Ah, but that reaction is in acidic aqueous solution, the reaction that produces HNO3 probably takes place in molten AN. So different conditions give
different reactions.
Eclectic - 17-11-2004 at 12:25
I did a patent search for "nitrous oxide ammonium nitrate chloride" Chloride must be excluded from molten ammonium nitrate production of
N2O because of the danger of EXPLOSIVE decomposition.
Just a warning that this is not a viable route to nitric acid production. If you want to actually try it on a small scale and it works, THEN
I'll believe it. I don't trust a single web page source with no references.
It will give some results
lucas - 19-11-2004 at 01:37
When distilling KNO3 and H2SO4. Both protons on the acid can be used. I have expeced this personally. I have otained yields of 1.2 moles HNO3 per mole
H2SO4.
The mixture is extremely hot, but the product is surprisingly good quality. The HNO3 boils so quickly on formation that it has little time to
decompose. All distillation was done under reduced pressure provided by a water jet vacuum pump. Measurement of the distillation mixture at this
stage of distillation showed the temperature to exceed the 250 degC range of my thermometer.
The prolem in starting with these ingredients will be getting them molten.
the melting point of sodium bisulphate is 186 degC (1). KNO3 melts at 334 degC (1). Once the bisulphite has melted, the reaction will start. The
mixture would need to be very intamately mixed before use. I would recommend use of .2 or so moles nitrate per mole bisulphite. Use of sodium
nitrate would be more desirable that potassium nitrate in this case, to maintain lower melting point ingredients and products.
Dont expect more than a small yield!
I dont think it's a usefull method!
(1) chemfinder.com
[Edited on 20-11-2004 by lucas]
Mumbles - 21-11-2004 at 12:08
If you are distilling under reduced pressure why(and how) in the hell are you distilling at 250C? IMO, there is no way the nitric acid can handle
these temps. I've have it start to decomposed to NOx with just a water bath, and sand baths below 150C. The point it of the vacuum is to be
able to distill at lower temps, like say 60C for example.
I have my doubts about this proceedure too. The temperature thing is what I'm worried about. Although, I suppose if it formed and cooled
quickly in the air it may be alright. One can remove dissolved NOx from nitric acid anyway.
unionised - 21-11-2004 at 14:15
I guess he's heating it that hot to get the reaction to work. Since it's the temperature of the mixture rather than the vapour (at least if
I read his post correctly) the vacuum might not affect it so greatly.
Distilling HNO3 under vacuum at 60C is a good way to ensure that the decomposition is minimised, it just isn't hot enough to get this reaction to
take place.
lucas - 25-1-2005 at 06:12
Yes, that is a massive temperature to distill under vacuum. The reason for the heat is that the reaction needs to be forced to occur. HSO4- + NO3-
--> SO4 + HNO3 is not a favoured reaction energetically. These temperatures are of course not used if sulphuric acid is present in significant
quantities. When the nitric acid is formed at this temperature, with lowered pressure, it immediately boils and is then condensed quickly after so it
has little time to decompose. This process yields yellow nitric acid and orange nitric oxide is profoundly present in the heated flask, but it does
produce usefull acid. Make sure you have a good supply of water to the condenser and don't overheat your thermometer. You may have to plug your
thermometer socket with a glass rod or stopper, depending on your still. This acid is fine for some uses, like nitric esters or can be cleaned up if
it's required for more fussy processes, like an RDX synth
[Edited on 25-1-2005 by lucas]
Costs
MadHatter - 25-1-2005 at 08:11
What about the costs of producing nitric acid via the bisulphate route ? If you live in an
area where sulphuric acid is unavailable or prohibitively expensive then you may not
have a choice. I would be concerned about yields of nitric acid and the decomposition
at those elevated temperatures.
The last bucket of sodium bisulphate I bought at the pool supply store cost me $15 for
10 LBS. I paid slightly less than that for my last 1/2 gallon of Rooto sulphuric acid.
I believe it's Polverone who can back me up on the prices for that fantastic acid.
frogfot - 25-1-2005 at 09:18
Wanned to remind that bisulphite decomposes on heating to water and SO3.. the latter can contaminate final product with H2SO4. Therefore, second
distilation would be recommended.
Btw, formed sulphate salt have a very high mp.. so closer to the end of distilling, reaction mix will become solid, leading to incompleate reaction..
maby it would be worth to find some inert solvent with high bp (>200*C)?
Theoretic - 25-1-2005 at 13:53
"Chloride must be excluded from molten ammonium nitrate production of N2O because of the danger of EXPLOSIVE decomposition."
Sorry for late reply, but...
It's all very dodgy, this "danger of explosive decomposition" line! AN is very hard to detonate, did anyone EVER have an N2O
preparation blow up in their face? At best (worst), you would have a fast deflagration-like decomposition when the critical 300 C threshold for =>
N2 + O2 decomposition is reached.
kclo4 - 12-3-2005 at 20:16
Would molten NH4HSO4 work instead with NH4NO3 to produce HNO3?
neutrino - 12-3-2005 at 20:50
It could work if you kept the temperature low. At higher temperatures, the ammonium sulfate would decompose, liberating ammonia.
kclo4 - 12-3-2005 at 20:56
NH4HSO4 ammonium bisulfate
neutrino - 13-3-2005 at 05:36
I was referring to these reactions:
NH<sub>4</sub>NO<sub>3</sub> + NH<sub>4</sub>HSO<sub>4</sub> -> HNO<sub>3</sub> +
(NH<sub>4</sub>
<sub>2</sub>SO<sub>4</sub>
(NH<sub>4</sub><sub>2</sub>SO<sub>4</sub>
-heat-> NH<sub>4</sub>HSO<sub>4</sub> + NH<sub>3</sub>
NH<sub>3</sub> + HNO<sub>3</sub> -> NH<sub>4</sub>NO<sub>3</sub>
12AX7 - 13-3-2005 at 06:34
Ah, kinda like:
CaCO3 + heat > CaO + CO2
CaO + H2O > Ca(OH)2
Na2CO3 + Ca(OH)2 > 2NaOH + CaCO3
repeat etc.
kclo4 - 13-3-2005 at 08:41
my bad