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Drhaines
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[*] posted on 8-6-2013 at 09:37
Question on HNO3


I've been doing some nitrations, and in utilizing the H2SO4/Nitrate salt method for the nitrating mix, I couldn't help but have the following question...

If you mix H2SO4 and, per say, NH4NO3, in stoichiometric amounts, resulting in the reaction of
H2SO4+NH4NO3--->NH4HSO4+HNO3, wouldn't the resultant supernatant mix be pure Nitric Acid(l) and Ammonium bisulfate(s)?
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[*] posted on 8-6-2013 at 10:35


No, not really. The nitric acid will still have significant amounts of ammonium bisulfate in it. The only way I know of to get the pure nitric is by distillation.



As below, so above.
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cyanureeves
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[*] posted on 8-6-2013 at 10:37


dont know but when i do it i get a mushy leftover instead of a solid and adding sodium hydroxide to the by-product will not release ammonia.oops! i thought you meant distilling.

[Edited on 6-8-2013 by cyanureeves]
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papaya
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[*] posted on 8-6-2013 at 12:44


Doesn't AN completely dissolve in H2SO4?
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cyanureeves
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[*] posted on 8-6-2013 at 14:29


AN does dissolve in sulfuric when distilling but it never left a solid behind as when i distill sodium or potassium nitrate with sulfuric.poorman's nitric acid is what original poster is talking about and it will never be uncontaminated.poorman's nitric made from sodium or potassium nitrate will always be contaminated too.i have done all three and it does dissolve silver and other metals but is highly contaminated and weaker than pure nitric acid.there is a video on you tube that shows nitrate salts in water added to sulfuric and then chilled until the salts settle on the bottom. the liquid remaining is called nitric acid but is highly contaminated. i dont think chilling can keep sodium or potassium sulfate from mixing with nitric acid if it is making contact with it.
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papaya
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[*] posted on 8-6-2013 at 14:50


With Ca(NO3)2 that might work.
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[*] posted on 8-6-2013 at 20:37


What about Ba(NO3)2 + H2SO4 --> 2HNO3 + BaSO4? Barium Sulfate is highly insoluble.



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[*] posted on 8-6-2013 at 20:49


The solubility of barium nitrate itself is quite low
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[*] posted on 8-6-2013 at 20:56


However calcium nitrate is quite soluble so it's possible that you could make, say 70% nitric acid by mixing H2SO4 and Ca(NO3)2 solution and filtering through a glass fretted Büchner funnel. I would be surprised if this hasn't been done already.
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[*] posted on 8-6-2013 at 23:08


Quote: Originally posted by DrSchnufflez  
However calcium nitrate is quite soluble so it's possible that you could make, say 70% nitric acid by mixing H2SO4 and Ca(NO3)2 solution and filtering through a glass fretted Büchner funnel. I would be surprised if this hasn't been done already.

The nitric acid would still be contaminated, since calcium sulfate becomes more soluble in acidic solution. This would probably still be a good way to separate out nitric acid for many purposes, however.
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Antiswat
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[*] posted on 9-6-2013 at 02:37


i know a guy who does it with Ba(NO3)2
he uses a ULTRA fine glass filter for it and vacuum
he says itworks so i will take his word for it

another more interesting part of this is that you could potentially do this with also Ca(NO3)2 as others mentioned, but instead of vacuum glass filters and more shiny glass you could put on some gloves, put the acid sludge into a piece of cloth and squeeze it
it should work
if this wouldnt work perhaps a piece of nitrated cloth would work?
the HNO3 would just nitrate the cloth when you squeeze it, i suppose you need alot of H2SO4 in it to actually start eating up the cloth
this could be combined with Ba(NO3)2 followed by decanting after several days
keep in mind 99% HNO3 is an actual dessicant and will fall in concentration if you use a huge container for it or store it for too long




~25 drops = 1mL @dH2O viscocity - STP
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[*] posted on 9-6-2013 at 03:32


Quote:
The only way I know of to get the pure nitric is by distillation.

But this is something many seem to find problematic, for whatever reason ─ I mean, an ordinary ground glass distillation apparatus is something anyone with an interest in chemistry should have as a basic piece of equipment . . .




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[*] posted on 9-6-2013 at 04:30


Please use the search facility, there are numerous threads on this subject. No point in reinventing the wheel over and over again.


[Edited on 9-6-2013 by blogfast25]




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platedish29
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[*] posted on 9-6-2013 at 17:17


Quote: Originally posted by Antiswat  
put the acid sludge into a piece of cloth and squeeze it
it should work

Try heated lead nitrate with sulphate, or silver with chloride... Both ways you get miserable results
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[*] posted on 11-6-2013 at 03:55


I started a thread a while ago on using Oxalic acid (available as H2C2O4 or H2C2O4.2H2O) in place of H2SO4 in preparations.

So use any nitrate just about (the general insolubility of the oxalate salts is the advantage, but check as some of them are soluble in acidic conditions) in a not too concentrated aqueous solution of your nitrate, and filter out the oxalate salt. For example:

2 KNO3 + H2C2O4 --> K2C2O4 (s) + 2 HNO3

However, avoid using an excess of H2C2O4 as the HNO3, under some conditions, apparently can be reduced to form vapors of Nitrous acid (see "The Chemical Gazette", Volume 12, pages 12 to 13 at http://books.google.com/books?id=JwcAAAAAMAAJ&pg=PA13&am... ).

Note, no distillation of some problematic acids is required or advised. The avoidance of producing a highly concentrated product is to avoid 'accidents' associated with preparing/working with too concentrated versions of many acids (yes, the problem is not too diluted and/or contaminated, but could arise from too pure and strong!).

Otherwise, you may be personally be made aware of some of the very nasty behaviors of some of these acids. For example, H2SO4 itself prepared this route, upon heating to concentrate (generally a bad idea) will reach a point where it self ejects from the test tube! Even a highly experienced chemist may be surprised by these concentrated acid properties, so in general, do not go there.

Oxalic acid will also change its role to a reducing agent, which may be part of the behavior issue discussed above. For example, in applying H2C2O4 in excess to a chlorate solution, it will first form HClO3 and, per its reducing properties, then explosive ClO2 gas from the excess H2C2O4.

Be also aware that H2C2O4 is poisonous, so do not insert your fingers in it, to see how it feels, or worst taste, or breathe in the dry dust. For example, here is a comment at a website (http://www.squidoo.com/use-of-oxalic-acid-in-the-shop-and-ar... ) that highlights oxalic acid many uses, but also warns, to quote: "Although oxalic acid is safer to handle than nitric, sulfuric, or hydrofluoric acids, oxalic acid is still fairly strong and quite poisonous so proper safety gear is important". See also MSDS at http://www.google.com/url?sa=t&rct=j&q=msds%20oxalic... ).

Finally, it is most likely not a low cost alternative to H2SO4 and its shelf life if not kept cool, is not as long either.

----------------------

As a side note in the case of preparing HNO3, since a RT solution path may work with H2C2O4, one could use NaNO2, if that is what you have cheaply and readily available:

2 NaNO2 + H2C2O4 --> Na2C2O4 (s) + 2 HNO2

Note, per Wikipedia to quote:

"In anything other than very dilute, cold solutions, nitrous acid rapidly decomposes into nitrogen dioxide, nitric oxide, and water:

2 HNO2 → NO2 + NO + H2O "

Also:

"In warm or concentrated solutions, the overall reaction amounts to production of nitric acid, water, and nitric oxide:

3 HNO2 → HNO3 + 2 NO + H2O "

So, a possible preparation route (I have not performed it myself), at RT create a strong solution of HNO2 using NaNO2 and Oxalic acid (but avoid excess H2C2O4 in this case), and then warm in a fumehood or outside to form HNO3. One could also capture and aerate the fumes to increase yield.


[Edited on 11-6-2013 by AJKOER]
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[*] posted on 11-6-2013 at 06:56


Quote: Originally posted by platedish29  
Quote: Originally posted by Antiswat  
put the acid sludge into a piece of cloth and squeeze it
it should work

Try heated lead nitrate with sulphate, or silver with chloride... Both ways you get miserable results


not sure if i understand what you mean.. Pb(NO3)2 decomposes into NO2 and lead oxides
so that wouldnt be useful for HNO3

lead nitrate is hard to come by and silver... mhm yeah.. thats if not only hard to come by also very expensive!!
it is doable but it would be extremely expensive as you would always loose some silver in the refining process
calcium nitrate is very easy to get in bulk quanities, and barium nitrate is too at fireworks suppliers




~25 drops = 1mL @dH2O viscocity - STP
Truth is ever growing - but without context theres barely any such.

https://en.wikipedia.org/wiki/Solubility_table
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AJKOER
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[*] posted on 11-6-2013 at 07:07


Quote: Originally posted by AndersHoveland  
Quote: Originally posted by DrSchnufflez  
However calcium nitrate is quite soluble so it's possible that you could make, say 70% nitric acid by mixing H2SO4 and Ca(NO3)2 solution and filtering through a glass fretted Büchner funnel. I would be surprised if this hasn't been done already.

The nitric acid would still be contaminated, since calcium sulfate becomes more soluble in acidic solution. This would probably still be a good way to separate out nitric acid for many purposes, however.


It is interesting to note that upon adding CaCl2 to a solution of MgSO4, it just forms a color of very dilute milk with no visible precipitate of CaSO4:

CaCl2 (aq) + MgSO4 (aq) <---> MgCl2 (aq) + CaSO4 (aq)
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[*] posted on 11-6-2013 at 07:33


What about phosphoric acid, could it drive HNO3 from the salt?
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[*] posted on 11-6-2013 at 08:04


Phosphoric acid is a weak acid, so probably would not be very effective is displacing nitric acid from the salt. Still, I think it would work with distillation. And I have also read of reactions where 70% nitric acid in anhydrous phosphoric acid is comparable to a mix of nitric and concentrated sulfuric acids.

But I certainly would not expected calcium phosphate to be able to precipitate out of a mix of Ca(NO3)2 and H3PO4.
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[*] posted on 11-6-2013 at 08:10


Quote: Originally posted by AndersHoveland  
Phosphoric acid is a weak acid, so probably would not be very effective is displacing nitric acid from the salt. Still, I think it would work with distillation. And I have also read of reactions where 70% nitric acid in anhydrous phosphoric acid is comparable to a mix of nitric and concentrated sulfuric acids.

But I certainly would not expected calcium phosphate to be able to precipitate out of a mix of Ca(NO3)2 and H3PO4.


I guess sulfuric and oxalic acid are the best way to go without distillation then. By the way, if someone wants a good route to get oxalic acid, it can be done by heating copper over a bunsen burner until it oxidizes and then putting hot into ethylene glycol. Procedure is on youtube, but I am not sure how they extracted it out of ethylene glycol. Maybe it's not even soluble in it?
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[*] posted on 11-6-2013 at 08:41


A method from my head for concentrating HNO3 by distillation:
If I take some dry phospate, say Na3PO4 and add to it diluted HNO3 in less than stochiometric quatity, this will displace weaker acid and I'll end up with NaNO3, H3PO4 and some excess Na3PO4 in water solution, right? Then if I start to distill this mixture, isn't it that first the water will boil out, then a dry H3PO4/polyphosphoric acids will displace HNO3 (more concentrated, because nearly all water is gone in the first stage) from NaNO3, which one can condense and finally Na3PO4 will be recovered at the end for further use.
In short - use phosphate instead of conc. sulfuric acid. Please point out flaws as I can't believe my head could generate something realistic at all :D .

[Edited on 11-6-2013 by papaya]
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[*] posted on 11-6-2013 at 09:22


Unfortunately this does not work. Na3PO4 is quite strongly basic and if you start from dilute HNO3, then you'll end up with a mix of NaNO3 (neutral), NaH2PO4 (slightly acidic) and Na2HPO4 (slightly alkaline) and all of these will remain back as a solid. When all water has boiled away, you'll have a mix of these three solids and you'll not recover any acid from this anymore.



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papaya
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[*] posted on 11-6-2013 at 09:47


You mean regardless of the HNO3 quantity taken (may be it was a mistake to say that I take less nitric than stochiometric) it'll not end up in a free H3PO4 and NaNO3? Also I meant that solid must react in the melt to give HNO3.
How I came to this idea is from another experiment: I had a waste solution of NaHSO4+some orgranics in it (maybe also Na2SO4, phosphates are not out) so I thought I can dry it by heating to obtain some impure bisulfate instead of throwing all out (can be used as an 'acid' somewhere). Thing is when I heated that solution and evaporated water down it made a bubbling melt (I learned that NaHSO4*H2O melts at low temp), which started to darken with a time (organics turning into carbon). I was disappointed with that carbon and redissolved solids in water, filtered and started all again, but this time I decided to add some 5-10ml of 50% HNO3 from the beginning to oxidize all organics during heating instead of carbonizing. This worked pretty well and I obtained nearly white stuff, however I was surprised when all water was already gone and the melt started to bubble - only then the flask started to give HNO3 vapors in a very intensive way, I thought it all gone much before! After some time it stopped to bublle and a clear melt obtained, which I believe is NaHSO4 with some contaminants. However bisulfate forms hydrate, so it cannot be used for this purpose, but are you sure phosphate is also out?
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