Sciencemadness Discussion Board

Concentrating Nitric acid

DubaiAmateurRocketry - 10-9-2013 at 11:08

Hi,

Because i need a high nitric acid concentration, preferbly over 95%, for something, I want to know how to concentrate the acid in easiest way.

First, I made my nitric acid by blowing NO2 gas bubbles into cold water. For around half an hour until the color becomes yellow, so i guess it has at least around 10% ?

Just trying to figure out how to concentrate it to >95% ?

Can i use salt that absorb water to become hydrated in order to suck the water out ? such as Lithium perchlorate ? or Copper sulfate ?

I would want to use Lithium Perchlorate because it has a low solubility in water when it is hydrated so yeah. And i have a lot of it any ways, and i can reuse it simply by dehydrating it on an oven so i think it might be convienient ?

Will that work ?

and how do i know how much of it shall i add when i am not sure of the concentrating of my nitric acid ?

How do i know if my filter paper will react with high % HNO3 ?

Metacelsus - 10-9-2013 at 11:55

First, titrate it to determine the concentration. Assuming it is low, I would convert it to a nitrate salt with a stoichiometric amount of alkali, evaporate the water, and then mix the salt with sulfuric acid and distill off concentrated nitric acid, preferably under vacuum. This distillation process has been heavily covered under other threads.

Adding a salt such as lithium perchlorate will just dissolve the salt in the nitric acid.

AndersHoveland - 10-9-2013 at 17:29

If one has 90% conc. nitric acid, and wants to concentrate it further, it can first be saturated with nitrogen dioxide and then a dry stream of ozone passed through. This can bring the concentration up to anhydrous, or even further. This route is most practical for already high concentrations, where one only needs to raise the concentration by a little more.

Variscite - 10-9-2013 at 18:01

Do you have a reference for that, Anders? Ive honestly never heard of that.

Pulverulescent - 11-9-2013 at 01:17

Anders is correct, but the method is obviously barely practical!
Distill 90% HNO3 from an equal volume of conc. H2SO4 to bring it to >95% . . .

Ascaridole - 11-9-2013 at 14:47

You may want to bubble your NO2 into hydrogen peroxide, will help your yield.
As for concentrating the acid I'm assuming you probably don't have access to concentrated Sulfuric acid?
If you do just carefully add the acid to the nitrate salt and distill off nearly anhydrous or possibly better nitric acid. Just be careful, your working with two very strong acids that can kill you.

If you don't have access to the concentrated sulfuric acid your best bet is to use distillation to get your 10% up to about 70% after that the azeotrope won't break easily without a strong dehydrating agent like sulfuric acid. You could try a modified method of Anders method by bubbling more NO2 into the 70% acid and blowing ozone through it. It will be a slow and expensive process but possibly doable.

Last but not least if you can find sodium bisulfate NaHSO you can mix it in a 1:1 molar ration with your nitrate salt and distill it to get nearly anhydrous nitric acid. It will be heavily contaminated with NO2 as the high heat needed will decompose a fair amount of your nitrate. An alternative is to use an aqueous mixture of the salts and distill off the 70% nitric acid azeotrope.

Either way making nitric acid and getting it to >95% is not a cheap or easy task. Even the stuff I use in the lab is around 70% (15M).

malford - 11-9-2013 at 18:41

Quote: Originally posted by AndersHoveland  
...it can first be saturated with nitrogen dioxide and then a dry stream of ozone passed through.


Could someone elaborate on this?

I am wondering, specifically, what is mean by saturating it with nitrogen dioxide.

Passing a dry stream of ozone through it could be accomplished with a setup such as this one.

Ascaridole - 11-9-2013 at 19:08

Essentially NO2 is dissolved in the solution by bubbling and ozone will react with it to form NO3. NO3 reacts with more NO2 to form N2O5 that gets hydrolized to HNO3. At least that's what I can remember.

malford - 11-9-2013 at 19:33

Pulvurulescent said that would be barely practical. Why?

What increases the solvation of nitrogen dioxide in the nitric acid? How about warming the nitric acid, small bubbles for high surface area, long travel distance for increased exposure?

Ascaridole - 11-9-2013 at 20:11

"Barely Practical" as in on an industrial scale it would be worth all the effort to setup the all the gas generators and stuff. As you said yourself to increase efficiency a fritted gas dispersion tube and a long travel distance are needed for the process to be practical. On a small scale such a setup would probably cost more than going out and paying for concentrated nitric acid. As for heating that would do more harm than good as nitric acid undergoes thermal decomposition.

I believe NO2 is more soluble in nitric acid than in water because NO2 gets hydrolized in water to nitric acid and NO (if bubbled in H2O2 no NO is produced increasing efficiency). As more and more HNO3 is made NO2 begins to exist in equilibrium with the acid. At least thats what I can figure out.

Fantasma4500 - 12-9-2013 at 09:08

i might be mistaken... but LiClO4 + HNO3 > HClO4 + LiNO3
HClO4 = you dont want that (:

whats the purpose of the high conc. HNO3 anyways? many nitrations are very possible with HNO3 below the azeotrope, otherwise just get some nitrate salt, that will give you anhydrous HNO3, but this isnt always preferred for nitrations as it tends to get hard to stir, especially if the ppt. is solid aswell

softbeard - 12-9-2013 at 14:46

It is also possible to use magnesium nitrate to concentrate nitric acid. I am fuzzy on the details because this is coming from memory of a book on the manufacture of nitric acid I read a long time ago, but I believe magnesium nitrate was used commercially to break the nitric acid/water azeotrope. Sorry I don\t have refs. and partial pressure tables.

ElectroWin - 12-9-2013 at 15:20

Quote: Originally posted by softbeard  
It is also possible to use magnesium nitrate to concentrate nitric acid. I am fuzzy on the details because this is coming from memory of a book on the manufacture of nitric acid I read a long time ago, but I believe magnesium nitrate was used commercially to break the nitric acid/water azeotrope. Sorry I don't have refs. and partial pressure tables.


i assume this works similarly to the way you can concentrate acetic acid by adding anhydrous CaCl2 and then distill?

i expect magnesium nitrate has a great affinity for water, therefore during distillation lets the HNO3 boil off first. So use anhydrous Mg(NO3)2.

Ascaridole - 12-9-2013 at 18:24

The magnesium nitrate method might work industrially but on a small scale I have a feeling it's another barely practical approach. Just as trying to dehydrate magnesium chloride results in formation of the oxide and HCl without careful controll of temperature the same is true for the nitrate. The only serious way an anhydrous or close to anhydrous magnesium nitrate could possibly be obtained would be through azeotropic drying via crystallization from an anhydrous ethanol solution or possibly drying at low pressure.

The truth be told nitric acid above the 68%-70% mark is hard even for professional chemists to obtain. Perhaps this is one synthesis we should accept as being difficult and dangerous without a professional setup and the proper reagents.

softbeard - 12-9-2013 at 18:59

From what I recall, the magnesium nitrate was made in-situ by adding magnesium oxide to the weak nitric acid. It would break the nitric acid/water azeotrope and hold on to some of the water while concentrated nitric distilled off. The leftover magnesium nitrate solution was then concentrated by boiling off some of the water and re-used. I agree, the method is probably trickier than it sounds since the hydrate ,Mg(NO3)2.6H2O ,does indeed decompose before it fully dehydrates by heat. I think the dehydration of the magnesium nitratewas performed under reduced pressure
Probably best to stick with sulfuric acid for the azeotrope-breaking.


[Edited on 13-9-2013 by softbeard]

DubaiAmateurRocketry - 13-9-2013 at 03:17

Quote: Originally posted by Antiswat  
i might be mistaken... but LiClO4 + HNO3 > HClO4 + LiNO3
HClO4 = you dont want that (:

whats the purpose of the high conc. HNO3 anyways? many nitrations are very possible with HNO3 below the azeotrope, otherwise just get some nitrate salt, that will give you anhydrous HNO3, but this isnt always preferred for nitrations as it tends to get hard to stir, especially if the ppt. is solid aswell


Umm, what about anhydrous Copper Sulfate ?

and are u sure nitric acid can do that ?!

[Edited on 13-9-2013 by DubaiAmateurRocketry]

Fantasma4500 - 13-9-2013 at 03:48

anhydrous CuSO4 as in for removing water from low concentration HNO3, right?
that should work, but gypsum should work aswell..

CuSO4 anh. can take per mole, up to 5 moles of water (5 x 16 mL / g H2O) thats 80 mL
but you need to keep in mind that you might need to squeeze the now wet CuSO4 to get more HNO3 out of it..
but however man... theres another approach..

get H2SO4, then boil it until it starts to fume, this should potentially work even at 70% to drag a larger amount of water onto the H2SO4, again this isnt something im fully sure about, you can keep on boiling it to 300*C, when its at 300*C and doesnt boil anymore you have around 98% H2SO4, which will surely work to remove the water
it can thereafter be added to boil off water and procedure can be repeated over and over with small H2SO4 loss during the concentrating from 70% to 98%
keep in mind SO3 is extremely nasty stuff if you decide to do this

otherwise theres yet against another possibility.. to get from 1 to 70% you could potentially use CaCl2 as dessicant, or silica gel, this would drag the water out of the HNO3, and you can remove dessicant, boil water off it so its dry and add to dessicant bag and repeat, and then use H2SO4 to get it to 99%

softbeard - 13-9-2013 at 08:56

Instead of copper sulfate, you might as well try anhydrous magnesium sulfate as your dessicant for the nitric acid. Probably it'll act more or less like magnesium nitrate (which it will partially form anyways when you add it to nitric acid).
Real cheap too.
Of course you have to distill off the nitric acid from a mix like this.

[Edited on 13-9-2013 by softbeard]

Ascaridole - 13-9-2013 at 18:16

Well if you have a ton of magnesium just reflux your 70% with magnesium shavings. You will loose some of your acid but it will generate your magnesium nitrate anhydrously and not produce a water in the process. You may not need to reflux it as the magnesium is pretty reactive but who knows, might passivate in the strong acid as more water is removed.

AndersHoveland - 13-9-2013 at 21:32

Nitric acid is also soluble in methylene chloride, so if you have a mixture of sodium nitrate and concentrated sulfuric acid, the nitric acid will be drawn into the methylene chloride layer.

Pulverulescent - 14-9-2013 at 01:12

The method has been used to produce anhydrous acid from 68%HNO3!
I've not seen mention of nitrate salts in connection with this extraction . . .

MrHomeScientist - 16-9-2013 at 06:11

Quote: Originally posted by malford  
What increases the solvation of nitrogen dioxide in the nitric acid? How about warming the nitric acid, small bubbles for high surface area, long travel distance for increased exposure?


Yes for the latter two, no for the first. Remember that gases dissolve better in cold solution, so warming would be counterproductive.

Also, it seems that a lot of people are recommending dessicants to concentrate your weak acid into stronger acid. This won't work. Dessicants are used to remove the last traces of water from an already nearly dry product - they can't pull out a gallon of water from 10% acid solution. (Unless you wanted to use 100 pounds of dessicant over a few months of drying time, I suppose)
If your acid is close to maximum concentration, then an appropriate dessicant (read: non-reactive with or soluble in your substance) can be used to bump it up a bit more.

Pulverulescent - 17-9-2013 at 06:09

Quote:
Remember that gases dissolve better in cold solution, so warming would be counterproductive.

Remember too ─ oxidation rate in soln. is temp. dependant.

Sunil Sharma - 17-9-2013 at 08:20

Since nitric acid makes azeotrope with water having boiling point 120 deg C and composition 68% nitric acid and 32% water... in this case when you distill 10% nitric acid solution you will first get the fraction of water and distillation residue will be be around 50/50 mixture of water and nitric acid... add 98% sulphuric acid to residue in equal quantity and again distill you will get aroud 95% nitric acid as distillate.....

AndersHoveland - 17-9-2013 at 12:45

Quote: Originally posted by malford  
Pulvurulescent said that would be barely practical. Why?

It is barely practical because it takes a large volume of ozone gas to make a very tiny volume of nitric acid. A typical ozone generator, whether produced by chemical means or electric discharge, is not very efficient, and typically the stream of oxygen coming out only contains 2-4% ozone. Add to this the fact that this gas stream has to be completely free from moisture, because chances are it contains a significant portion of water vapor relative to the small concentration of ozone present, and the gas will have to flow through the nitric acid for a long time. This route is not practical for making any more than very small amounts of nitric acid. The only advantage is that it is not constrained by a maximum concentration.

Hydrogen peroxide can also oxidize NO2 to nitric acid, but the hydrogen peroxide typically contains water. It also leaves behind water after it oxidizes the NO2. Even assuming you had pure H2O2, the maximum concentration of HNO3 theoretically obtainable would be only 77.7%.

In contrast, ozone only leaves behind O2 after it oxidizes the NO2. You do not need ozone to oxidize aqueous solutions of NO2, it can be oxidized to nitric acid by regular air, but the maximum possible concentration achievable by oxidation with air is only around 70-77%. That is because O2 can not actually directly oxidize NO2, but rather relies on an equilibrium that exists in aqueous solutions of NO2, an equilibrium that becomes less favorable as the concentration increases.

If you are trying to obtain nitric acid through distillation much past the azeotropic concentration, reduced pressure needs to be used, otherwise the heat will irreversibly decompose a portion of the nitric acid vapor.

[Edited on 17-9-2013 by AndersHoveland]

Pulverulescent - 18-9-2013 at 01:19

Concentration of HNO3 by electrolysis!

PHILOU Zrealone - 18-9-2013 at 03:02

Quote: Originally posted by Antiswat  
anhydrous CuSO4 as in for removing water from low concentration HNO3, right?
that should work, but gypsum should work aswell..

CuSO4 anh. can take per mole, up to 5 moles of water (5 x 16 mL / g H2O) thats 80 mL
but you need to keep in mind that you might need to squeeze the now wet CuSO4 to get more HNO3 out of it..
but however man...

Actually, it would be slightly better 5x18g of water = 92g or ml water ;)

The process with MgO (or CaO) or Mg(NO3)2 (or Ca(NO3)2) and filtration and/or distillation is interesting because:
1°) dehydratant is recyclable by heating and too much a heating causes not too much arm because MeO (Me = Mg or Ca) is also compatible with the process except it uses some of the HNO3.
2°) dehydratant at a certain nitrate content will cristallize and allow easier filtration due to precipitation by common anionic effect.
3°) no interferences of external anions...like sulfate, chloride

malford - 19-9-2013 at 11:21

Would not distilling again and using a high number of theoretical plates in the distilling column be the most efficient and possibly effective method of reaching very high concentration?

Pulverulescent - 20-9-2013 at 05:04

In distilling from H2SO4/NaNO3 mixtures water finds its way into the distillate from decomp. of HNO3!
Distillation under good vacuum sidesteps this . . .

malford - 20-9-2013 at 08:50

I would be distilling from nitric acid and sulfuric acid under a vacuum of 18 inches of mercury while being stirred to prevent uneven heating.

Further, with a sufficient distilling column, one theoretically would not even need to distill again. I'm going to try a hemple column that is 25mm x 300mm filled with roughly 500 0.25-inch glass balls. By carefully controlling the temperature in the column using insulation, it seems possible that the product making its way over to the condenser could be nearly pure.

Pulverulescent - 21-9-2013 at 03:29

Fractionation isn't necessary or desirable . . .
The ordinary water-cooled condenser will work fine for non-vacuum distillation, while for vacuum work extra cooling is advisable!
And, anyway, there are very few cases where anhydrous acid is essential!

malford - 21-9-2013 at 05:31

I don't see how fractionation is undesirable. Would not each evaporation and condensation cycle leave behind more of the less volatile compounds?

Anhydrous acid is, in fact, my goal.

Pulverulescent - 21-9-2013 at 07:29

Quote:
I don't see how fractionation is undesirable.

Nitric acid behaves as a constant-boiling fraction in such distillations, regardless of its water-content.

malford - 21-9-2013 at 18:59

If the proper amount of sulfuric acid is introduced to the nitric acid, it seems to me that there theoretically would be no water left to form the azeotrope.

What would be left, it appears to me, would be:

H3O+
HSO4-
SO42-

[Edited on 22-9-2013 by malford]

Bot0nist - 21-9-2013 at 20:47

The amounts may have less to do with the affinity sulfuric acid has for water, than the thermodynamic favoribility of a small fraction of water to form an azeotrope with HNO<sub>3</sub>. I am sure much more complicated explinations of vapor phase interactions are required, but im at a loss...

I do know that if you apply a halfway decent vacuum to a
standard distillation apparatus (no fractionating column needed) you will get nitic acid likely beyond suitable for your needs. Even without a vacuum, very strong RFNA is easily obtained with even just 1 to 2 distillations starting from an anhydrous nitrate salt and some ~98% sulfur acid. Bubble dry air through to remove dissolved NO<sub>x</sub> and what more could you need?

[Edited on 22-9-2013 by Bot0nist]

malford - 23-9-2013 at 08:35

I agree that extremely pure nitric I will probably not need. But, I'm pursuing the science of it.

As such, I'm curious about using a distilling column. There is a large difference between a distilling column not being needed, as in more than required for typical applications, and the same being counterproductive. Is it in fact counterproductive?

Pulverulescent - 24-9-2013 at 01:45

Well, the greatly extended vapour-path requires strong heating to drive the distillation ─ so yes . . .

malford - 24-9-2013 at 04:47

Ah, so you're saying it is counterproductive because of the potential for decomposition of the nitric acid?

Well, if a vacuum is used and the column is insulated, I suppose we could circumnavigate those worries. In such a case, would not the column become an asset and improve the purity of the product?

My main concern is, what exactly are the products of a nitric acid, sulfuric acid mixture, what are their boiling points, are any of them azeotropic with nitric acid? Suppose I start with 70% HNO3 and 98% H2SO4.

Pulverulescent - 24-9-2013 at 04:54

Quote:
Suppose I start with 70% HNO3 and 98% H2SO4.

Yes do, the vapour will be of constant comp. with or without a column . . .

malford - 24-9-2013 at 12:38

When you say vapor, what compounds are you referring to? It seems there are lots of products of a nitric acid (70%)/sulfuric acid (98%) reaction. I am unable to find the properties of all of the products and thus cannot determine whether a column would provide a benefit. I am hoping to determine this scientifically rather than anecdotally.

Bot0nist - 24-9-2013 at 12:45

The do some experiments, with and without a column, and analyze the obtained fractions...

The condensate is what your after, after all. It will be nitric acid, and a very small amount of water, either carried over from the reactants, or from the decomposition of the acid.

malford - 24-9-2013 at 12:51

If you would like to visit me and help determine the existence and precise amount of every product compound, I would be very grateful. Otherwise, I do not have the knowledge or equipment to do so.

Edit:

http://www.jacobs.com/uploadedfiles/Products/Chemetics/Nitri...

In this document, it is said in the third paragraph that adding sulfuric acid (50% weight) "completely" eliminates the azeotrope. It is said in the fifth paragraph that a distilling column is then used to produce up to 99% pure nitric acid.

They may be using the column to separate leftover acids from the pure nitric acid rather than to gradually purify the nitric acid. However, their language "further concentrated" referring to the increase from 80% to 99% nitric acid in the upper portion of the column seems to imply that the column employs fractionation.

[Edited on 24-9-2013 by malford]

testimento - 19-10-2013 at 11:50

I found the following data (similar to malford post):

- Nitric acid of 68% azeotrope can be eliminated with min. 50% H2SO4 at 1:1 volume (1L az. HNO3 + 1L min. 50% H2SO4) by fractional distilling the nitric acid at 83C to obtain concentrations up to and over 99%, and then recycling the H2SO4 by boiling off the water to get high conc. again

- Magnesium nitrate anhydrate (dec. 330C) can be used in similar manner mixed by equivalent weight (1kg&1L) to eliminate the azeotrope. Because of the low decomp point, vacuum fractionation and water bath heating is preferred method

Questions:

- Could HNO3 be dehydrated with CaSO4?
- Could H2SO4 be dehydrated with CaSO4?

Off-topic for nitric acid:

-HNO3 can be manufactured by calcium or magnesium nitrate, at 650C or 330C, and an oxide is leftover and NO2 gas is released, which forms HNO3 upon water.
-CaO can be reacted to form CaOH with water
-MgO can be reacted with NaOH, in where Mg(OH)2 is formed and precipitated as solid(practically insoluble), and Na will instantly scavenge new OH group from water, and no NaOH is consumed in the transfer process
-Mg(OH)2 can be reacted with ammonium nitrate, in which metathesis will form Mg(NO3)2 and ammonia, released as gas upon heating - CaOH applies exactly the same manner in this phase
-The purpose of this transfer process is to eliminate completely the consumption of sulfuric acid to obtain HNO3

-NO2 forms nitric acid in water, and one NO group is released in air. Therefore, one should use long plastic or steel cylinder, for ex. 200mm diameter and 1500mm high to scrub the NO2, because NO is slightly heavier than air, and it stays in the cylinder and re-oxidizes to NO2, which condenses back to the water, and by this method, the recovery rate of NO2 gas can be increased over 95%. This produces max. 68% HNO3 without pressurizing the scrubber, which can be concentrated as mentioned before.

AJKOER - 14-11-2013 at 15:23

Here is an alternate perspective which may be justified on safety concerns alone (very concentrated acids are dangerous and, at times, their behavior is unexpected).

The idea is not so much focusing on concentrating HNO3, but on raising the 'activity level' of the acid in the particular application. Adding Ca(NO3)2 or Mg(NO3)2 may accomplish this.

A cited example occurs in hydrometallury, where one can significantly raise the 'activity level' of dilute HCl by adding MgCl2 (alternatively, to a lesser extent using NaCl). This is an important tool in this field where one is trying to leach out minerals from ores efficiently and cheaply.

Source: See Hydrometallurgy in Extraction Processes, Vol I, by C. K. Gupta and T. K. Mukherjee, page 15 at <del>http://books.google.com/books?id=F7p7W1rykpwC&pg=PA203&lpg=PA203&dq=FeCl2+%2B+O2+%3D+FeO(OH)+%2B+FeCl3&source=bl&ots=fi WLs05y8f&sig=mi-pV94woVj7JABKBB zLZcqbEwM&hl=en&sa=X&ei=ZQgfUeq7BIi50AGynoDoBQ&sqi=2&ved=0CFAQ6AEwBA#v=snippet&q=Magnesium%20chloride%20MgCl2&f=false</ del> http://books.google.com/books?id=F7p7W1rykpwC&pg=PA15 .

where the author cites data confirming that a 2M HCl in 3M CaCl2 or MgCl2 (or FeCl3) behaves like 7M HCl.


[Edited on 14-11-2013 by AJKOER]

<!-- bfesser_edit_tag -->[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: del; fixed link?]

[Edited on 20.11.13 by bfesser]

blogfast25 - 15-11-2013 at 05:43

Quote: Originally posted by AJKOER  

where the author cites data confirming that a 2M HCl in 3M CaCl2 or MgCl2 (or FeCl3) behaves like 7M HCl.


[Edited on 14-11-2013 by AJKOER]


That link get truncated when someone else uses it.

Can you cite these data?

AJKOER - 19-11-2013 at 07:08

OK, I actually check this messy link and it does work. The 'trick' is to copy and paste it. However, one must copy the entire link text including the last space and period.

To be honest, there is precious little more data then I presented. However, there is more research (some advanced) on the topic of activity level. I would also not consider myself an exceptional source on this topic, but I have taken a course in Numerical Analysis and have source books. This latter discipline being applicable as the actual calculations of activity levels in practice could involve the solution of a system of non-linear equations. Apparently, there are companies that consult for industrial chemical producers by solving these numerical problems.

blogfast25 - 20-11-2013 at 08:42

It works for you (but not me), probably because it's cookie based.

Personally I'm sceptical about acid activity increasing by adding a salt of the conjugated base to it. Even HCl + H2O < === > H3O+ + Cl- is an equilibrium (albeit with a very high Ka), so stuffing the solution with extra Cl- should push the equilibrium a bit to the left.

Subjectively I've also experienced that acids are less reactive to metals in solutions of high ionic strength (all other things being equal). I have no clear empirical evidence for it though.

[Edited on 20-11-2013 by blogfast25]

bfesser - 20-11-2013 at 09:12

AJKOER, spaces cannot be used in a URL. Your browser is likely auto-correcting your copy-pasted link from history. I replaced the link in your post; please let me know if it's not in alignment with your intent.

watson.fawkes - 20-11-2013 at 15:35

Quote: Originally posted by blogfast25  
It works for you (but not me), probably because it's cookie based.
It's Google Books, which is known to have issues crossing jurisdiction.

Practicaler - 18-4-2014 at 04:18

I have lots of 40% nitric acid can I simply mix it with 98%H2SO4 to distill more than 90% concentration HNO3

Zyklon-A - 18-4-2014 at 08:08

Yes, but it would be more practical to first distill it by itself to reach the azeotrope, then use sulfuric acid to reach near anhydrous.

Etanol - 23-8-2020 at 12:30

I found in the literature table of dryers activity for nitric acid: Al(NO3)2>Fe(NO3)3>Mg(NO3)2>Ni(NO3)2>H2SO4>LiNO3>Al2(SO4)3>Zn(NO3)2>Ca(NO3)2>H3PO4>Cd(NO3)2>NaNO3>KNO3.
Has anyone tried to distill nitric acid with LiNO3? LiNO3 is inferior to Mg(NO3)2, but it does not decompose when dehydrated. I could not dehydrate Mg(NO3)2 2H2O to anhydrous salt without obtaining NOx. The distillation with Ca(NO3)2 gives only 90% nitric acid.

[Edited on 23-8-2020 by Etanol]

morganbw - 23-8-2020 at 13:59

Quote: Originally posted by Zyklon-A  
Yes, but it would be more practical to first distill it by itself to reach the azeotrope, then use sulfuric acid to reach near anhydrous.


This is the correct answer.