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Author: Subject: golfpro's Nitric Acid for Beginners
golfpro
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[*] posted on 24-8-2013 at 06:07
Nitric Acid yeilds from Retort?


Hey for those using retort, how much concentrated nitric acid are you getting per distillation with what amount of sulfuric and KNO3? I don't have a problem with the purity, but I use 250ml sulfuric acid and then the proper amount of KNO3 I forgot which, and then only get 50ml of Nitric acid??? Am I losing my acid in the way of it just not condensing and fuming off instead? I don't cool the tube on the retort but the acid drops into a ice cold beaker. I use an oil bath at least 150* C


my retort only holds 500ml, and I don't like 300ml sulfuric yeilding only 50ml Nitric... This is with the correct amount of KNO3.
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[*] posted on 24-8-2013 at 06:32


Try using a cold round-bottomed flask. This will provide a larger cold surface.



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[*] posted on 24-8-2013 at 06:48


Quote:
I don't cool the tube on the retort but the acid drops into a ice cold beaker.

On distillation, the neck of the retort heats quickly until the glass reaches the vapour temperature and the vapour then largely escapes!
You could try pushing a cooled condenser onto the neck using teflon tape to minimise leaks . . .
Otherwise, just use small quantities of reactants ─ fiddly, but you'll lose less product!
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[*] posted on 24-8-2013 at 07:33


i got a 500ml retort and never used it because of that same thing but using less reactants is what i would do also.i even thought about attaching an aluminum tube to the neck and sticking out through a hole in my tool shed wall.i would probably have to only distill in the winter though and hope the aluminum passivates.
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[*] posted on 24-8-2013 at 13:10


The distillation would work much better in the winter, but is there any way to cool the neck like wrapping a baggy of crushed ice and water and taping it around the neck?

This would cause moitsture to condense on the neck and maybe even drip water into the beaker which is supposed to get high concentrated nitric acid and then defeat the purpose.

I bet I'd get at least twice the amount of acid distilling in 15*F weather compared to 85*F Until then I've got to find a way to cool the neck. Maybe a cloth dipped in ice cold water and wrap that around..
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[*] posted on 24-8-2013 at 15:01


<strong>golfpro</strong>, I'd appreciate it if you replied to your existing threads on very similar (or identical) topics, rather than starting new ones. I don't know how things are run in EM, but out here, your new topics are excessive.

[edit] I've merged three of your topics on nitric acid so far, and know there is at least one more. I recommend using the search function before posting, and posting replies in appropriate threads rather than starting new ones. <img src="../scipics/_warn.png" /> <em>If you open any more beginnings-level topics on nitric acid, I will likely lock or remove them.</em> <img src="../scipics/_warn.png" />

[2nd edit] Oh, to hell with it! I merged all five of them&mdash;I don't care if it's a little messy; this is low-level stuff.

[Edited on 25.8.13 by bfesser]




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[*] posted on 27-8-2013 at 10:22


Quote: Originally posted by Pulverulescent  

The cooled vapour condenses, dissolving NO2 and lemon-coloured acid drips from the condenser!
A small quantity of NO2 escapes the condenser, depending on cooling-water temp. but the distillate contains most of the NO2 produced!


The method I normally use to get rid of the NOx is bubbling dry air through the nitric acid while heating it to about 50 degrees C.
Would bubbling O3 do the job quicker? I'm considering building an O3-generator as a hobby project so if it is usable to eliminate the NOx that would be a nice application for it.

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[*] posted on 28-8-2013 at 02:22


I hardly expect the O3 from a standard O3-generator to be of any influence. Of course, the O3 will oxidize the NO2 (in combination with water) to HNO3, but usual O3-generators only have a few tenths of percent at most of O3. The rest is air (or oxygen, if you use an oxygen supply). I have never seen pure O3-gas and I think that making this is very difficult.



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[*] posted on 26-9-2013 at 16:09
maximizing distillation yields from a retort


Has anyone found a way to get maximum distillation yields from a retort? I have a 500ml retort and from 200ml sulfuric acid and 160g KNO3, I get 50ml of Nitric Acid. I am looking for a way to get 100ml Nitric acid per 200ml sulfuric acid.

Maybe the neck slightly more angled down? I know that on a cold day I'd get more HNO3, and with less material in the retort, more nitric acid can be attained, but this only means it's more efficient to use less each time.

I was thinking of a way to cool the neck, maybe with a bag of ice?

I figure there are people on here who used a retort and have found little things for more efficient distillation.
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[*] posted on 26-9-2013 at 16:18


Ive never used a retort but the best way I can think of to increase your yield without using an actual condenser would be to stick the end as far as possible into another flask submerged in an ice bath so the entire flask acts like a condenser as well as trying to cool the neck as much as possible.

http://www.youtube.com/watch?v=S8rtyRnZZMU

Theres a good visual of what I'm trying to describe but definitely don't use copper and steel for making nitric acid :D.

You might also try cooling the flask with dry ice, that will cause the nitric acid to freeze but it shouldnt crack your glassware as it does not expand like water.

[Edited on 27-9-2013 by PeeWee2000]




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[*] posted on 26-9-2013 at 16:31


Ah the retort. I own and periodically use a retort from time to time. Unfortunately, they are just not that perfect for good distillations...but not entirely without purpose. The idea is to try and keep the heat at just the minimum you need to distill off whatever it is you are trying to isolate. With the retort, you don't have the advantage of a chilled condenser, so try wrapping the retort in a cold wet rag (and replace the rag with fresh cold wet rags as they heat up). Also, make sure your collection flask is sitting in an ice bath.

Ideally, you can eventually invest in a simple distillation setup. A round bottom flask, a distillation head, a liebig condensor, and a take off. You can probably pull this off for just under a hundred dollars on Ebay. You can even use a simple automotive siphon to circulate water through the condenser. And after you obtain this setup, it's fairly simple to upgrade to fractional or vacuum distillation by simply buying a fractionating column or vacuum adaptor. Then you'll be distilling like the big boys!

[Edited on 27-9-2013 by MichiganMadScientist]
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[*] posted on 26-9-2013 at 16:33


Another thing I think of (I haven't used retort either) would be to wrap a towel soaked in freezing water (Maybe with some dry ice wrapped inside) and wrap that around the retorts condensing end.
Depending on how long the distillation takes you may have to have a couple of towels ready.
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OHp, michiganmadscientist beat me to it!

[Edited on 27-9-2013 by Dariusrussell]
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[*] posted on 26-9-2013 at 17:36


Sell it to some sucker whos into "alchemy", and use the $ to buy a proper, ground glass, distillation apperatus...



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[*] posted on 29-9-2013 at 13:13
NO2 removal from anhydrous nitric acid


I know it's been talked about and I've searched but this thread can just be deleted in 2 days to save space.

If I have Pure Nitric Acid aside from NO2 contamination, how can I remove this without losing the concentration? Some places I see pasing CO2 gasses through drives the NO2 out, sometimes I read just regular air... I've heard of adding small amounts of Urea...

I sat for a few minutes with a pipette making bubbles and saw no color change, is it a much longer process than that? How would urea react w/ the nitric acid to drive out NO2 while keeping it very pure? Or is CO2 needed?
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[*] posted on 29-9-2013 at 19:03


NO2 is extremely soluble in anhydrous concentrated sulfuric acid, it is taken up and reacts. The reaction can be described by the equilibrium:

2 NO2 + 3 H+ <==> NO+ + NO2+ + H3O+

The nitrosyl and nitronium ions actually exist in the form of nitrosylsulfuric acid and nitronium hydrogen sulfate. Nitrosylsulfuric acid is surprisingly stable, it can even precipitate out as a solid in only moderately concentrated solutions of sulfuric acid (with plenty of water present), though excess water will lead to hydrolysis.

For lower concentrations - say only 90% - the solubility of NO2 drops suddenly because the species nitronium hydrogen sulfate can no longer exist in the presence of water. (However, mixtures of nitric oxide and nitrogen dioxide are easily taken up by more moderately concentrated solutions of sulfuric acid, because then just nitrosylsulfuric acid can form).

In anhydrous nitric acid, however, this uptake equilibrium is not so favorable because of all the nitrate ions present, but I would imagine it still plays a significant role in the solubility.

NO+ + NO3- --> 2 NO2


If you want to rid anhydrous nitric acid of a discoloring nitrogen dioxide impurity, the best route may be to pass a flow of dried ozone in, to oxidize the NO2. This should only be used to remove the last remnants of NO2 impurity that cannot be removed by other methods first used.

[Edited on 30-9-2013 by AndersHoveland]




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[*] posted on 29-9-2013 at 19:28


I have had success by passing dry air from a fish tank bubbler through a drying tube into a graduated cylinder of nitric acid. In the morning, the 100ml of acid was very clear...



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[*] posted on 29-9-2013 at 19:35


One thing to consider is that anhydrous nitric acid is very deliquescent, it will pull out traces of moisture from any gas being passed through it. If a large volume of gas is being passed in over an extended period of time, it may be important to ensure that the gas is free from moisture. This can be achieved by passing the gas through a column of finely pulverized baked calcium chloride.

CO2 seems like a good gas to use to help pass out and displace most of the NO2, if there is a lot of NO2 in there. Another possibility may be reduced pressure, though one suspects the acid fumes would probably be very corrosive on your vacuum pump.

[Edited on 30-9-2013 by AndersHoveland]
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[*] posted on 30-9-2013 at 04:28


How efficient is crystallization for HNO3 purification? The melting point of nitric acid is -42. Is slightly impure HNO3 (traces of water, NO2 et cetera) in the -40-s a viscous substance that is poorly crystalizable, or is it easily frozen leaving excess water and NO2 in the liquid?
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[*] posted on 30-9-2013 at 10:21


Ok, if I bubbled CO2 gas continuously for 10 minutes through would this do it? I know there are many variables, but I'd like to know if that would be a 10 hourish process or 10 minutes.
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[*] posted on 15-10-2013 at 18:55
68% concentration of a solution, by weight or volume?


If we have 1L of 68% Nitric acid, is there 680ml of pure HNO3 and the other 320ml is water? Or is 68% of the Liter By MASS composed of pure HNO3 with the rest being water?


[Edited on 16-10-2013 by golfpro]
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[*] posted on 15-10-2013 at 19:03


My understanding is when nothing else is specified, the concentration is in "weight by weight" (w/w) basis. So when you have 1 lit of 68 % Nitric acid you actually have 1 * (density of 68% HNO3) * 0.68 kg of 100% HNO3 dissolved in 1 * (density of 68 % HNO3) * 0.32 kg of water.

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[*] posted on 15-10-2013 at 22:32
by weight


68% nitric acid means 68% of the weight is HNO3 and 32% of the weight is H2O. The density of HNO3 is going to be different (thicker) than the density of H2O, so the volume of HNO3 is going to be less than 680ml per liter of acid. You need the density of nitric acid to compute relative volumes.
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[*] posted on 15-10-2013 at 23:26


I'll back up the previous answers. A solution expressed in % is always a proportion of solute mass to total solution mass. 68% HNO3 means 0.68g HNO3 per 1.00g acid solution. The other 0.32g should be water in this case, but this won't be the case with every other solution out there (stabilizers, buffers, etc). I generally always convert % to M as a matter of personal preference, mass ratios can be misleading.
For example:
-Large molecules (especially if you get into organics) can have very high mass % concentrations and still be quite low in molar concentration.
-1L of solution is NOT equal to 1000g of solution. This is a characteristic of water at room temperature.
-You really shouldn't measure out a mass% solution by volume unless you've done the proper calculations (densities, Molarity, etc).

My two cents:P, sorry if I ranted, this insistence on % strength ratings has always been a pet peeve.
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[*] posted on 23-10-2013 at 17:47
NaNO3 instead of KNO3 for nitric acid distillation


Hi,

I have lost my source of KNO3 which I distilled from H2SO4 to produce HNO3, I may no longer be able to produce 99% Nitric acid unless I find something to replace the KNO3, so would NaNO3 work as a substitute here?

there is more Nitrate concentration w/ sodium nitrate than potassium nitrate..

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[Edited on 24-10-2013 by golfpro]
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[*] posted on 23-10-2013 at 18:09


NaNO3 will work just fine. Make sure to readjust the stoichiometry as needed. Ammonium nitrate from cold packs works as well.
(I wouldn't give up so quickly - even if all your OTC sources disappear, there are always online sources...)




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