Sciencemadness Discussion Board - Can I concentrate my 68% nitric acid in to 90%+?

Can I concentrate my 68% nitric acid in to 90%+?

LD5050 - 21-2-2017 at 17:18

How did you get 90% nitric acid I thought the highest concentration was 68% How d you break the azeotrope?

[Edited on 2-22-2017 by zts16]

Cryolite. - 21-2-2017 at 17:33

Quote: Originally posted by LD5050

How did you get 90% nitric acid I thought the highest concentration was 68% How d you break the azeotrope?

You just distill nitric acid by mixing a nitrate salt and sulfuric acid. If there is only a small amount of water in the acid and the salt, the nitric acid which comes off can only contain this little bit of water, and will therefore be stronger than azeotropic.

Can I concentrate my 68% nitric acid in to 90%+?

LD5050 - 21-2-2017 at 17:43

I made some azeotropic nitric acid by distilling a mixture of sulfuric acid, ammonium nitrate, and water. I know have 68% Nitric acid but I need a more concentrated form to oxidize Benzyl alcohol into Benzaldehyde. Is there any way to break the azeotropic nitric acid that I already have? I don't believe I am able to distill fuming nitric acid from H2SO4 and ammonium nitrate because ammonia gas produced will neutralize the nitric acid.

LD5050 - 21-2-2017 at 18:06

I only have ammonium nitrate tho and it produces ammonia gas and neutralizes the nitric acid so Is there any way to concentrate the 68 to 90%?

Lambda-Eyde - 21-2-2017 at 18:06

Yes, there is. You can distill it again with more concentrated sulfuric acid to retain the water. However, if you didn't know this or couldn't find this out on your own I don't think it's a very good idea. Fuming nitric acid is seriously nasty stuff even when you're not distilling it. The distillation process produces some nitrogen oxides as well.

Texium - 21-2-2017 at 19:03

Quote: Originally posted by LD5050

I only have ammonium nitrate tho and it produces ammonia gas and neutralizes the nitric acid so Is there any way to concentrate the 68 to 90%?

It shouldn't, when you have sulfuric acid in there. You'll be left with ammonium sulfate, which doesn't begin to decompose until well over 200ºC. Ammonium nitrate is fine, though potassium nitrate or calcium ammonium nitrate is preferable.

LD5050 - 21-2-2017 at 20:49

Quote: Originally posted by Lambda-Eyde

Yes, there is. You can distill it again with more concentrated sulfuric acid to retain the water. However, if you didn't know this or couldn't find this out on your own I don't think it's a very good idea. Fuming nitric acid is seriously nasty stuff even when you're not distilling it. The distillation process produces some nitrogen oxides as well.

I did a quick search on google and through other sources but couldn't find much other than starting from scratch which I didn't want to do, and also I only had ammonium nitrate and read that it cant be done because it neutralizes the acid. I figured I would probably be able to distill the 68% from more H2SO4 to obtain a higher concentration but i wasn't sure so I asked on here. Thanks tho, and I do know the dangers associated with concentrated acids. What sulfuric to nitric acid molar ratio should I prepare in the boiling flask when doing this? I'm sure its not super important to get it precise.

Texium - 21-2-2017 at 21:00

To be honest, you're overthinking it. Your best bet is to prepare sodium or potassium nitrate by adding your ammonium nitrate to a boiling solution of NaOH or KOH (or carbonate if you don't have a hydroxide available). Ammonia gas will be driven off, and you'll be left with a solution of the pure alkali nitrate. I'd recommend potassium over sodium because it is much easier to crystallize (if your boiling solution is fairly concentrated, you should be able to recover most of it just by cooling it).

LD5050 - 21-2-2017 at 22:08

Ya I actually just did this procedure, 200g ammonium nitrate and and 100g of sodium hydroxide. I didn't boil it tho except from it boiling on its own for a couple minutes after I poured in the water. After the reaction died down I put some more water in (around 200ml total) and mixed it around and then filtered which was a pain. I put it in the freezer to initiate crystalization but not much seemed to happen so I boiled off some of the water until I saw some precipitate and now i got it cooling on the counter. I see it crystalizing out but the yield seems pretty low.

chornedsnorkack - 21-2-2017 at 22:14

Can you end up with above 100 % nitric acid?

LD5050 - 21-2-2017 at 22:58

Quote: Originally posted by chornedsnorkack

Can you end up with above 100 % nitric acid?

How would you come up with nitric acid above 100%?...thats impossible...

Tsjerk - 22-2-2017 at 00:41

Nitric acid with its dehydrate N2O5 dissolved in it is referred to as being more than 100%

Herr Haber - 22-2-2017 at 04:03

STOP ADDING WATER !

Only add water to your round bottom flask AFTER the distillation to get rid of the sulphate salt.
Adding water when you are trying to get a concentrate is a bit counterproductive to say the least.

Fulmen - 22-2-2017 at 05:20

Quote: Originally posted by LD5050

What sulfuric to nitric acid molar ratio should I prepare in the boiling flask when doing this?

Equal volumes of 98% H2SO4 and 65% HNO3 is a good place to start.

There is also a process using magnesium nitrate instead of sulfuric acid. Never tried it but it could be easier to regenerate the magnesium nitrate.

chornedsnorkack - 22-2-2017 at 09:44

What´s better to concentrate nitric acid: H₂SO₄ or P₂O₅?

PHILOU Zrealone - 22-2-2017 at 15:16

Quote: Originally posted by chornedsnorkack

What´s better to concentrate nitric acid: H₂SO₄ or P₂O₅?

N2O5 is far better

Booze - 26-2-2017 at 13:25

Quote: Originally posted by LD5050

I made some azeotropic nitric acid by distilling a mixture of sulfuric acid, ammonium nitrate, and water. I know have 68% Nitric acid but I need a more concentrated form to oxidize Benzyl alcohol into Benzaldehyde. Is there any way to break the azeotropic nitric acid that I already have? I don't believe I am able to distill fuming nitric acid from H2SO4 and ammonium nitrate because ammonia gas produced will neutralize the nitric acid.

When you did this, did you distill nitrogen dioxide? I have been trying to do this for some time. Any tips?

ficolas - 26-2-2017 at 13:39

Quote: Originally posted by Booze

When you did this, did you distill nitrogen dioxide? I have been trying to do this for some time. Any tips?

Do you, for some weird reason want nitrógen dioxide, or do you want nitric acid? As somebody alredy told you, nitrogen dioxide is a product of the decomposition of nitrix acid, if you want nitric acid, you dont need to "distill nitrogen dioxide"

Fulmen - 26-2-2017 at 13:47

Quote: Originally posted by chornedsnorkack

What´s better to concentrate nitric acid: H₂SO₄ or P₂O₅?

Define "better".
P2O5 is a potent chemical, capable of dehydrating even stubborn chemicals/mixtures. But regenerating phosphoric acid back to P2O5 isn't possible. So unless you have good access to P2O5 and have a use for the phosphoric acid H2SO4 will be better as it can be re-concentrated by distillation.

chornedsnorkack - 26-2-2017 at 14:28

Just to remind: the pure compounds in the system are:
HNO₃ - boils at +83, freezes at -42
N₂O₅ - melts at +41, boils at +47
N₂O₄/NO₂ - boils at +22, freezes at -11.

LD5050 - 28-2-2017 at 03:27

Quote: Originally posted by Booze

Quote: Originally posted by LD5050

When you did this, did you distill nitrogen dioxide? I have been trying to do this for some time. Any tips?

Very little brown/red nitrogen dioxide was formed I also used NileRed method. After the first distillation the acid will be pretty weak especially if its cold. I added this to some copper with not much effect at all. I then proceeded with a fractional distillation and collected everything after 120c and ended up with azeotropic nitric acid which I got a VERY strong reaction when added to copper. The entire process is pretty straight forward I'm confused as to what you are having trouble with ( I saw your other post). I would suggest to fractionally distill the distillate after the first distillation.

Tsjerk - 28-2-2017 at 03:36

You don't need highly concentrated nitric acid to oxidize benzyl alcohol.

Amos - 28-2-2017 at 07:29

Tsjerk is correct; if you even read my original benzaldehyde thread you'd have seen all the papers that achieved even better results than mine using dilute nitric acid and a small amount of nitrite.

chornedsnorkack - 28-2-2017 at 10:11

Azeotropic, 68 % nitric acid is near an eutectic at about 71 %.

LD5050 - 28-2-2017 at 17:49

Quote: Originally posted by Amos

Tsjerk is correct; if you even read my original benzaldehyde thread you'd have seen all the papers that achieved even better results than mine using dilute nitric acid and a small amount of nitrite.

Amos I did in fact read your thread and I wanted to try the 10% nitric acid method but I didn't have any nitrite and without out it the yields aren't to great.

Tsjerk - 1-3-2017 at 07:50

As you need the nitrite in catalytic amounts you could reduce a bit of nitrate with lead which is very doable in small amounts.

AJKOER - 1-3-2017 at 17:00

Actually, rather than focusing on concentrating Nitric acid, instead try adding N2O (or, a nitrite, see http://cires1.colorado.edu/jimenez/AtmChem/CHEM-5151_S05_L7.... ) to even dilute HNO3 and treat with UV/strong sunlight in the presence of the substance to be nitrated. The targeted photolysis product is the nitrate radical, which while short lived, I suspect may be a stronger nitrating agent that your super strength HNO3! More precisely, NO3• can add to double bonds or engage in hydrogen abstraction (see page 4 at https://www3.nd.edu/~ndrlrcdc/Compilations/Ino/Ino.pdf ).

The underlying mechanics:

N2O + H2O + hv ---> N2 + OH- + •OH

•OH + HNO3 ⇆ H2O + NO3•

Rate constant for the forward reaction estimated as (8.6 ± 1.3) × E07 L mol-1 s-1

Reference: see "Equilibrium constant of the reaction .OH+HNO3←→H2O+NO3. in aqueous solution" by Poskrebyshev, G. A.; Neta, P.; Huie, R. E., published in Journal of Geophysical Research: Atmospheres, Volume 106, Issue D5, pp. 4995-5004, 03/2001. Link: http://adsabs.harvard.edu/abs/2001JGR...106.4995P

With respect to my possible reactivity claims, some specifics, see, for example, "Reactivity trends in reactions of the nitrate radical (NO3) with inorganic and organic cloud water constituents" by Hartmut Herrmann et al. To quote from the abstract:

"The rate coefficients of the reactions of NO3 with acetone (reaction 2), tetrahydrofuran (reaction 3), and hydrated formaldehyde (reaction 4) were determined at T = 298 K to be k2 = (4.4 ± 0.6)· 103 L mol−1 s−1, k3 = (2.1 ± 0.1 )· 107 L mol−1 s−1, and k4 = (7.8 ± 1.2)· 105 L mol−1s−1, respectively. For the reactions of NO3 with the alcohols methanol (reaction 5), ethanol (reaction 6 ), 1-propanol (reaction 7), 2-propanol (reaction 8), and 2-methyl-2-propanol (reaction 9 ) the following rate constants at T = 298 K were obtained: k5 = (5.1 ± 0.3)·105 L mol−1 s−1, k6 = (2.4 ±0.5)·106 L mol−1 s−1, k7 = (3.2 ± 0.1)· 106 L mol−1 s−1, k8 = (3.7 ± 0.9)·106 L mol−1 s−1, and k9 = (6.6 ± 0.7)·104 L mol−1 s−1, respectively."

Link: http://www.sciencedirect.com/science/article/pii/00167037949...

[Edited on 2-3-2017 by AJKOER]

LD5050 - 1-3-2017 at 23:29

Quote: Originally posted by Tsjerk

As you need the nitrite in catalytic amounts you could reduce a bit of nitrate with lead which is very doable in small amounts.

Tsjerk is there a write up or some more info on this that I can read up on? I would be interested in this if I have the materials.

Tsjerk - 2-3-2017 at 02:06

AJOEKER, as lovely theoretic your post always are, I wouldn't put your advice in practice. I wouldn't add N2O to any reaction that doesn't need N2O, especially in this case as it is a oxidation reaction and not a nitration. It would just give more benzoic acid.

LD5050, I'm on my phone and don't want to search, but have a look on this board, it is discussed in detail.

[Edited on 2-3-2017 by Tsjerk]

AJKOER - 2-3-2017 at 04:25

My suggestion was either N2O or nitrite (more light sensitive per my cited reference than dinitrogen oxide, and nitrates also, but much less photo active) under photolysis employing dilute HNO3. Interestingly, note the prior comment below by Amos:

Quote: Originally posted by Amos

Tsjerk is correct; if you even read my original benzaldehyde thread you'd have seen all the papers that achieved even better results than mine using dilute nitric acid and a small amount of nitrite.

Could it be that fluorescent lab lights were photo-catalysts for benzaldeyde in the presence of nitrite, or a trace transition metal impurity producing hydroxyl radicals in a Fenton-like fashion, or even the action of (or in combination with) air/O2 and light on any present transition metal ions?

Related research with organics, see http://www.publish.csiro.au/en/pdf/EN10004 .

[Edit] The general effect of radiation, light or micro wave irradiation or electric current on N2O I would described as:

N2O + e- ----> N2 + O-

which can be oxidizing, but in the presence of water:

O- + H2O = OH- + •OH (See reaction [14] at https://pdfs.semanticscholar.org/d696/b35956e38351dd2eae6706... )

where the equilibrium is far to the right except under highly alkaline conditions (or no water).

Also, my suggested trial use of N2O is due primarily to its low price and wide availability of 8 gram cartridges sold in major stores for the purpose of beverage infusion, which is also cool, but perhaps not healthy for sun bathers/drinkers as I once speculated on SM. Note, no credit card generally required, disclosure of mailing address, possible watch list, etc., which may be associated with various internet (or sometimes store) purchases of various chemicals (including NaNO2) by home chemists.

[Edited on 2-3-2017 by AJKOER]