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Author: Subject: Nitric Acid Synthesis
Magpie
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[*] posted on 9-8-2010 at 08:00


Your original acid is below the azeotropic concentration of 68% so the vapor condensed will be rich in water. When the distillate temperature reaches 120.5C your pot concentration is that of the azeotrope. Does this make sense?

See this excellent write-up:

http://www.chemguide.co.uk/physical/phaseeqia/nonideal.html

The last page has the phase diagrams for the water/HNO3 system. It's been very helpful to me.

[Edited on 9-8-2010 by Magpie]

[Edited on 9-8-2010 by Magpie]
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[*] posted on 9-8-2010 at 08:40


Yes! Perfect sense, and it makes life easier too. Boil down to azeotrope, then run some conc sulphuric in and distil off the nitric.
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[*] posted on 13-8-2010 at 05:20


Does anyone care to tell me what vacuum is their choice for clean distillation of greater than azeotropic nitric? What temperature seem a good cross between clean processing and easy condensation using water?

Will a fridge compressor pull this vacuum or do I need better?
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hissingnoise
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[*] posted on 13-8-2010 at 05:55


What's wrong with the strength of HNO3 distilled from KNO3/H2SO4 at normal pressure?
Its concentration is sufficient to nitrate pentaerythritol and hexamine, so I can't see where the problem is. . .

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[*] posted on 13-8-2010 at 06:11


I have glass setup doing just that at the moment. All the product is seriously red/brown. I was hoping to make WFNA directly as locally sulphuric is a problem so I don't want to waste the reagents making and discarding the brown NOx.

Also I can buy 35% nitric easily so it should be possible to concentrate it (yes azeotrope I know!). The winchester of Sulphuric was a chance ebay lot, with a long drive to get it. Even stores that sell sulphuric drain cleaner elsewhere in the UK don't have it near me so I'd rather work on hydroponic nitric (35%) which just arrives by van simply for the buying it.
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hissingnoise
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[*] posted on 13-8-2010 at 06:54


The dissolved oxides can be removed (oxidised) by blowing dry air or better, dry oxygen, through your acid.
The last traces may be slow to disappear and a trace of added urea will finish the job.

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uchiacon
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[*] posted on 21-9-2010 at 20:08


Hey so I just bought a new aspirator to replace my old one. The old one was metallic, and since I was concentrating nitric acid in the distillation, it really didn't last a long time. I thought that I would be able to trap enough of the acid vapours in a vapour trap, but no such luck. As you can see in the picture, the nice nozzle that used to supply the vacuum is pretty corroded.

I got a 2nd hand glass aspirator, with a bit of what I assume is a leaf in the nozzle chamber. It doesnt effect the vacuum though I dont think. Got it for a great 40NZ or about $28US.
It uses a lot less water than the metal one, and it has enough of a vacuum in the nozzle chamber to make the water boil inside the aspirator, which is at about 12 degrees.

In a sealed setup, it can boil water at 25 degrees, and it might do it at 20 degrees, but I don't really know that as I havent tried to boil it at 20 degrees. I'd assume there is a way you can calculate this into mm/hg or PSI or kPa, but I don't really have a clue. It's about 1/4 of an atmoshphere that my aspirator is pulling right? I think it's a lot better than my metal aspirator, which could only boil water at 35 degrees C, but it takes a bit longer to evacuate the glassware. Seems to pull a better vacuum than the metal one too.

Could someone point me towards a page that has a good rundown on pressure? I've had a look on wiki, but I just end up getting a bit confused, the formula
Pressure = exp(20.389-5132/temp C) gives me a number with 185 zeros behind the decimal point of torr.

What am I doing wrong here?
I used a calculator on the internet, and it said that water will boil at 25C at 23.8torr/mmhg. Is this right?




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[*] posted on 21-9-2010 at 21:01


Just use the sigma nomograph :

http://www.sigmaaldrich.com/sigma-aldrich/areas-of-interest/...

If waters normal bp is 100c, and you are getting it to boil at 25c, then your vacuum is 56mmhg.



[Edited on 22-9-2010 by Chainhit222]




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[*] posted on 22-9-2010 at 05:36
nitric acid


ive seen it done with mere glass coffee pot on B.B.Q. grill. with oil lamp tops used as condensing tubes and another coffee pot at opposite end.the more and wider the tubes the better it condenses. youll be surprised how much heat this glassware can tolerate.teflon tape was wound around the glass tubes and shoved into the mouth of the (mr.coffee)recepticles.bailing wire was wrapped around the smaller circumferance ends and used as flange rims to hold the glass set up tightly together. never saw the pressure build and push the apparatuses apart nor great amount of red fumes escaping but i do recall an annoying nose irritation upon standing close by.
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[*] posted on 26-9-2010 at 00:20


guys, I was just doing a distillation. But after distilling 150mls of 70% HNO3 with 300ml of 99% H2SO4, I got a yield of 20mls, at which point nitric stopped coming over the still head and down the condenser.

The 20mls I did get was 95-96% concentration, burnt a small crater in my hand when a small droplet landed on part of the skin that the glove didnt protect. I didn't quite realize the potency of this acid until that happened lol.

20mls for 150mls isn't good enough for my liking. Can you guys diagnose the problem?

I use an aspirator to pull a vacuum in the setup I use. I have a 40cm liebig condenser that I use. The nitric acid I was distilling was boiling and coming over at around 70 degrees, and then it stopped boiling and then nitric acid seemed to stop evaporating, but I think 100mls of nitric acid has left the original nitric/sulfuric acid mix. Has this 100mls of acid or water just been sucked out by the aspirator? I think my joints are pretty tight, and I used 1 wrap of teflon tape to make them more airtight.

I didn't cool the condensated acid with ice, just water that was about 16 degrees.

Here's a vid I made a while back with a similar setup, except the aspirator was bigger than the small glass one I have now.
www.youtube.com/watch?v=yXxSAQFo9ZE&feature=related

How should I go about reducing the amount of acid lost to the aspirator?




I also have a big 25L container of 70% nitric which I bought for $3/L, and I was hoping to use this up by concentrating it with the big 15L container I have of 99% sulfuric which I got at $2.75/L ;) ..... but would it just be better to distill nitric from scratch and have it concentrated that way? If it did it under vacuum, would it achieve a nice 90%+ conc of nitric?

[Edited on 04-07-09 by uchiacon]




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[*] posted on 26-9-2010 at 00:49


OK you should have 105ml but be satisfied at 100!

Firstly moderate the vacuum! While working at low pressure reduces the decomposition it also makes condensing the product much more difficult. If you want to use the aspirator then moderate the vacuum by adding a small air bleed into the vac line. Also cooling the condenser water may help. Where was the vapour condensing down the centre tube. It's best if the top quarter is vapour but the rest is condensing liquid, If the pressure is too low then the liquid in the receiver flask will evaporate too -Use an ice bath there too!
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[*] posted on 27-9-2010 at 22:24


ok, well I did my second distillation today.

Instead of 20mls 95%, I got 90mls 96%.

Instead of using a 2:1 sulfuric/nitric acid mix, I used a 1:1 acid mix of 200mls each. I used 70% nitric and 99% sulfuric acids.

This time I cooled the receiving flask with ice and water. My 60cm liebig condenser had a decent amount of 14 C water going through it.

The acid mix began boiling at 45C, and it was still boiling with large bumps infrequently as I went to turn it off and remove the vacuum. At this point, the temperature had risen to 60 degrees C.

I did not notice a large amount of nitric acid in my aspirator vacuum tube, but there was maybe a few mls. It was flattened out because of the vacuum inside it.,

I'm still not satisfied with this yield, should I have let it boil longer? Should I reduce the vacuum?

Any help will be appreciated... thanks




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[*] posted on 28-9-2010 at 08:57


At 45c with a 14C condenser you are challenging distillation. You could well use cooler water (it will come with winter!) or better try to get the main boil at round 60c by moderating the vacuum.
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[*] posted on 7-11-2010 at 01:16
Recycling H2SO4


Hello everyone,

Reading some of your very informal posts, I was thinking of how one can recycle the sulfuric acid used in nitric acid production.

My idea is to use NH4NO3 + conc. sulfuric acid in excess to produce the nitric by distillation (as normal) and then heat the ammonium sulfate + conc. sulfuric acid residue (at ~290 degrees). By that way, one can release ammonia and sulfuric acid by decomposition of ammonium sulfate; the sulfuric acid stays in the flask.
References: http://www.freepatentsonline.com/EP1444166.html

What do you think, is it worth trying?
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[*] posted on 7-11-2010 at 03:16


Ammonium sulphate does dissociate at high temperature but its components, rather than separating, recombine on cooling - and so it is hard to take seriously the patent you cited . . .
Na and K nitrates with H2SO4 give more concentrated HNO3 than NH4NO3/H2SO4 and are preferred for this reason!


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[*] posted on 7-11-2010 at 05:33


Thanks for input.
As I see it in the patent, the ammonium sulfate is decomposed inside concentrated sulfuric acid. So, the sulfuric acid from the decomposition, is mixed with the existing acid while ammonia boils off, thus effectively separating those two substances (at least, this is how it is supposed to work).
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[*] posted on 7-11-2010 at 06:21


The reaction mentioned does occur to some degree but at a rate far too slow for practical purposes.
Also, hot, conc. H2SO4 is a potent oxidiser which can attack NH4, forming N2, H2O and SO2.
NH4NO3 and (NH4)2SO4 in aqueous solution @>100*C evolve traces of NH3.
(NH4)2SO4 is reduced to the bisulphate and the nitrate becomes increasingly acidic but at an almost immeasurably slow rate.
BTW, the patent itself contains a peppering of misspellings, making it a bit unusual . . .

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[*] posted on 7-11-2010 at 08:23



Quote:

The reaction mentioned does occur to some degree but at a rate far too slow for practical purposes.


Yes, the possible low speed of the reaction is my major concern also. Also, probably there will be some contamination in the acid from the various decomposition products.
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[*] posted on 7-11-2010 at 09:43



Quote:

Na and K nitrates with H2SO4 give more concentrated HNO3 than NH4NO3/H2SO4 and are preferred for this reason!


Can ammonium nitrate give azeotropic (68%) HNO3? I plan to dissolve the required NH4NO3 in hot H2SO4/water solution. Also, in order to avoid decomposition of NH4NO3, I plan to distill the acid in vacuum.
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[*] posted on 7-11-2010 at 09:54


NaNO3 gives good yields of HNO3 @ normal pressure!
The NaHSO4 you're left with is more useful than (NH4)2SO4 . . .

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[*] posted on 7-11-2010 at 12:04



Quote:

The NaHSO4 you're left with is more useful than (NH4)2SO4 . . .


Why? Can I make sulfuric acid from that?
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[*] posted on 7-11-2010 at 12:50


Well yes, as it happens, you can - if you have the determination . . .

http://www.sciencemadness.org/talk/viewthread.php?tid=10217#...


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[*] posted on 7-11-2010 at 14:21


Some of the strategy known as UTFSE would have led to knowledge that it is the bisulfate of ammonia NH4HSO4, not the normal sulfate of ammonia (NH4)2SO4, which is the byproduct of HNO3 synthesis using NH4NO3 + H2SO4. Heating the ammonium bisulfate results in dehydration to ammonium pyrosulfate.

To get a normal sulfate byproduct from the reation of a nitrate and H2SO4, a nitrate must be selected which does not form an acid sulfate. Aluminum nitrate is one example. Such a reaction leads to a twice greater yield of HNO3, per mole of H2SO4, since both of the acid hydrogens of the H2SO4 are then used to form HNO3, rather than the usual one acid hydrogen which remains with an acid sulfate byproduct.
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[*] posted on 7-11-2010 at 23:52


Thanks for correcting me.
What is pyrosulfate? Is it the same as ammonium persulfate, (NH4)2S2O8? In wikipedia it says that it decomposes at 120 °C. In what substances does it decompose?
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[*] posted on 8-11-2010 at 03:34


Quote:
Thanks for correcting me.

And me!
I was under the impression that NH4HSO4 in H2SO4 readily converted to the normal sulphate @~60*C and that the corresponding reaction with NaHSO4 was slow @ >125*C!
Oh well . . .
Anyway, persulphates are salts of peroxydisulphuric acid H2S2O8.
Ammonium persulphate decomposition produces a range of products which includes ozone and NO2.
Sodium pyrosulphate (Na2S2O7) thermally decomposes to Na2SO4 + SO3.







[Edited on 8-11-2010 by hissingnoise]
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