Sciencemadness Discussion Board - Our Beloved Nitric Acid

Our Beloved Nitric Acid

DetaDude - 5-10-2009 at 17:06

Since alot of what we do involves the use of this valuable commodity (especially the energetic materials crowd) I thought that a new post on this topic might be in order. I know this topic has been talked to death in other posts & threads, however I would like to hear from some our more advanced and experienced citizen chemists (no I'm not one of them).

The idea behind this post/thread is to amass some plain vanilla type info for our more basic and novice people.

I am 68 yrs old and have been at pyro/chemistry/explosives since age 11. That said I would like to generate on this post a wealth of knowledge on the topic of just nitric acid.

What works and what does not, what is better vacuum or non vacuum distillation, column or no column, ratio of H2SO4 to HNO3 when distilling if it is even used, alternative methods to concentrating/purifying etc., etc., etc.

When I need a high concentration of HNO3 I use 70% HNO3 and 96% H2SO4 under 29,5 in. Hg with a column and get excellent results, to wit 98% HNO3. But this is just my method how many other ways are there to boost the percentage of HNO3, and are there any heavy expenses involved?

Please try to keep it plain vanilla, at least freshman high school chemistry, remember I would like this to be for the lay person.

Thanks ahead of time to all those that respond

Rich_Insane - 6-10-2009 at 18:49

You and I live in the same area.

I am having trouble just making 70% HNO3. Mostly because H2SO4 is in Ace hardware and I can't drive, and I don't have a legitimate distillation system.

I must say, there is already a stickied thread about HNO3 production.

ammonium isocyanate - 6-10-2009 at 19:09

Depending on quantity, you don't really need any fancy equipment to make nitric acid.

I made about 20mL fuming nitric acid using the following apparatus:

A 50mL beaker was placed inside a 250mL beaker, and the to the outside edge the 250mL beaker was added sodium nitrate (anhydrous) and 98% sulfuric acid (1 to 1 molar ratio). A funnel with the stem broken off (the stem was already broken-- I didn't break it off just for this) was placed over the 250mL beaker (which it just covered, leaving a small opening where the spout was) such that the funnel was centered over the 50mL beaker. Several small strips of Teflon tape were balled up and used to plug the funnel so that it would hold water (which worked suprisingly well- no water ever came through it). The funnel was then filled with crushed ice. The apparatus was then placed on wire gauze and heated over a burner until distillation ceased.

I didn't do quantitative analysis of the nitric acid, but it was light yellow, presumably due to dissolved NO2. The nitric acid reacted vigorously with silver to produce AgNO3 and brown fumes of NOx.

Rich_Insane - 6-10-2009 at 19:23

I see how that'd work. Now I have to get some 50 ml beakers....

Well, I was thinking of producing 100 ml of HNO3 at a time, as opposed to buying 2.5 L of HNO3 at ridiculous prices. I only really need it for nitrations ATM.

halogenstruck - 10-10-2009 at 11:54

it is not concentrated,i suppose.
con. HNO3 does not react vigorously with Ag as silver nitrate is sparingly soluble in concentrated aid and makes a protective layer which protect it against furthur reaction.
i know it for 60%-50%
but maybe it`s not correct for over 80% as i did not used it over Ag before.
30% and lower corrode silver good

blogfast25 - 10-10-2009 at 12:26

Quote: Originally posted by halogenstruck

it is not concentrated,i suppose.
con. HNO3 does not react vigorously with Ag as silver nitrate is sparingly soluble in concentrated aid and makes a protective layer which protect it against furthur reaction.
i know it for 60%-50%
but maybe it`s not correct for over 80% as i did not used it over Ag before.
30% and lower corrode silver good

It would help if ammonium isocyanate could quantify a bit by weighing the nitrate, sulfuric acid and the obtained nitric acid using his distillation method. This could give rough idea of yields and concentration of the nitric. Otherwise, titrate with NaOH...

halogenstruck - 11-10-2009 at 12:50

yes u r right,method is genius and adding some details about the yield could be a good idea.
also i would add up to my last post here,i think if Ag is not completely pure like containing 10% Cu which is normal,it may dissolve much vigorously.
i had 99%HNO3 and it was yellow too.it is possibly because of dissolved NO2/NO gases.
although i think concentrated solutions of NO2/HNO3 tends to be green.when exposed to air,turns to colorless due to NO2 loss

hissingnoise - 12-10-2009 at 04:58

N2O3, dissolved in HNO3 colours the solution green. . .
NO2 is always amber in colour.
The dimer N2O4 is colourless!

ammonium isocyanate - 12-10-2009 at 17:54

I no longer have the old sample of nitric acid which I obtained using the method I described, so I rebuilt my apparatus today and ran a small scale test. If absolutely necessary, I will perform a titration on the sample I produced to verify concentration.

Apparatus:
-Outer beaker was 150mL
-Inner beaker was 50mL
-Funnel, 75mm, with a broken stem, was plugged using wads of teflon tape and placed resting on the lip of the outer beaker with the remainder of the stem located directly above the inner beaker.
-Reactants were poured into the gap between the inner and outer beakers.
-Funnel was filled with ice water.

Procedure:
4.84g of anhydrous sodium nitrate was added to the outer beaker, followed by 6.22g of sulfuric acid (an 11% molar excess assuming end product is NaHSO4). The plugged funnel was filled near the brim with ice water, and the whole apparatus was heated over a burner on low heat until distillation from outer beaker slowed significantly and the beaker became red-brown from nitrogen dioxide. The apparatus was set aside to cool, and the funnel was removed. 2.0mL of distillate was isolated, density 1.48g/mL (theory, 100% nitric acid, 4.52g/mL, 70% nitric acid, 4.42g/mL), which in contact with air fumed to produce a white, highly irritating, generally nasty mist.

Conclusion:
This method requires very little in the way of glassware, and appears to produce fuming nitric acid (based on density and qualitative observations). It could also likely be scaled up with ease to several times the scale used in this experiment. Furthermore, it is likely that one could use less sulfuric acid, substituting greater heat in order to drive the reaction to Na2SO4. However, even with these modifications, it is unlikely that a large amount of nitric acid could be produced by this method.

Apparatus:

halogenstruck - 12-10-2009 at 18:35

it fumes because it vaporizes to air but with the steam in air makes its azeotrope with water with higher boiling point and lower volatility.as a result it condenses again as small droplets and makes white fume.it fumes only in humid air.sometimes it makes fume because of NH3 in atmosphere
good job,u made a concentrated HNO3 with a simple apparatus

ammonium isocyanate - 12-10-2009 at 18:39

Yes I figured the fuming was as such. The air where I live is almost always humid, so as such I always deal with such things. The fuming really was impressive though, much more than 37% hydrochloric acid.

Engager - 19-10-2009 at 17:27

Acid produced by conc. H2SO4 / NaNO3 reaction is not anhydrous it's just about 95-97% in general and contains some nitrogen oxides witch are unacceptable in some nitrations since NOx promote oxidative action of acid. To make pure colorless HNO3 from acid mentioned above, one must measure it's density and dehydrate it using calculated ammount of P2O5, remove NOx by passing stream of dry air or CO2 and finaly distill it in vacuum.

[Edited on 20-10-2009 by Engager]

halogenstruck - 20-10-2009 at 08:56

a drop of H2O2 maybe makes completely white although a little bit diluted

a_bab - 20-10-2009 at 10:30

My 5 cents on concentrating HNO3: instead of using the expensive and hard to find DCM (dichloromethane) for extraction, just use plain old magnesium nitrate anhydrous. It'll dty the nitric acid very well, leaving it white pure.

No farting around with distillation setups and such.

hissingnoise - 20-10-2009 at 11:52

AFAIK, Mg(NO3)2 is usually supplied as a hydrate and trying to dehydrate it by heating is fruitless as this salt is decomposed by heat.

Theophrastus_2 - 20-10-2009 at 16:05

Well, the simplest method (and the one I happen to use) simply involves taking almost any nitrate salt (potassium, sodium, etc.) and placing it within an erlenmeyer flask. You then add some 33% hydrochloric acid to the mix, stirring so that all of your nitrate salt dissolves in the acid- water mixture. You should then prepare the apparatus; a second erlenmeyer flask, containing water. As well, ready a holed stoppers (coated in teflon or something of the sort, as nitric acid can eat through certain types of rubber) as well as some tubing (check if the material is resistant to nitric acid), which will vent the fumes produced in your first erlenmeyer flask, to the second. Once your apparatus is set up, add copper metal to the HCl/NO3, and seal it off with the holed stopper, with the connected tubing/ piping. The vessel will quickly begin to produce NO2 gas, which gives HNO3 when bubbled through water. Based on your stoichiometry, you can control the concentration of your final solution. (granted your yield will never be 100%)

And I agree about having some sort of publication on making nitric and hydroiodic acid and so forth, for more inexperienced chemists (I don't deny, in some respects I remain in that category). Hell, if it weren't for piling schoolwork, I might do it!

Theophrastus_2 - 20-10-2009 at 16:06

Damn, that shouldn't have been most efficient; not simplest.

hissingnoise - 21-10-2009 at 03:34

Your 'method' is more waste of time and nitrate than method. . .
If you want HNO3 in any quantity or concentration get yourself a proper distillation apparatus and forget that other rubbish!

DetaDude - 6-12-2009 at 20:32

Has anyone tried concentrating nitric with a reflux set up, if so how about a little feed back on how it was done, and the outcome, equipment used, was vacuum involved, did you also use a column, etc.

crazyboy - 6-12-2009 at 21:07

Quote: Originally posted by DetaDude

Has anyone tried concentrating nitric with a reflux set up, if so how about a little feed back on how it was done, and the outcome, equipment used, was vacuum involved, did you also use a column, etc.

Why would you try to concentrate anything with a reflux setup?

DetaDude - 15-12-2009 at 16:39

I guess I should have made myself clearer. The intention was, has anyone tried using a reflux condenser in the distillation of HNO3, with or without a column or vacuum.

entropy51 - 15-12-2009 at 18:34

Still not clear.

A reflux condenser is not a particular kind of condenser (Liebig, West, Allihin, Friedrichs) but rather a condenser arranged for reflux, to return condensate to the boiler. Any of those types of condensers can be used for reflux.

No one can understand how you can distill this way, for good reason. Try describing exactly what you have in mind if you want a meaningful answer.

hissingnoise - 16-12-2009 at 08:10

If you mean fractional distillation it won't work because there's no way round the azeotrope. . .
Its boiling temperature is always 120*C at normal pressure!

[Edited on 16-12-2009 by hissingnoise]

DetaDude - 17-12-2009 at 22:08

entropy51:

Sorry but you are amiss with your statement about a reflux condenser. There is a condenser that is made just for reflux operations and it is very similar to a Graham type condenser except the cooling fluid flows through the coil instead of the jacket. This is the item that I had reference to in my earlier post.

I use this item atop a column when distilling certain liquids to increase condensate purity, I also use a vacuum (29.0 in Hg).

I was just trying to get input from others that may have had some experience with this type of setup.

bbartlog - 18-12-2009 at 06:03

Regardless of whether there is such a thing as a 'reflux condenser', entropy's basic point still stands: if you're *refluxing*, then in theory everything is dripping back into the vessel, so no purification or concentration is going to occur. If on the other hand you're assuming that some volatile fraction (water or w/e) is actually escaping, then what you're doing is more like fractional distillation. I can imagine some cases where a sort of hybrid setup would be useful, say if you wanted to reflux ethanol while allowing aldehyde vapors to escape and condense elsewhere, but that's not really what we're dealing with here...

entropy51 - 18-12-2009 at 06:28

Quote:

There is a condenser that is made just for reflux operations and it is very similar to a Graham type condenser except the cooling fluid flows through the coil instead of the jacket

Sorry Dude, but that condenser is often used to distill solvents like ether, in which case it is not refluxing, and thus not "just for reflux". It is sometimes called a "reversed Graham condenser".

It sounds as if you are using it as a fractionating column? Or what? It would be nice if you were able to describe what you are doing in a comprehensible fashion. If you could do that you might get a useful answer, which you have not so far, have you? GIGO.

hissingnoise - 18-12-2009 at 15:05

Deta, fractional distillation works on concentrations above ~68% but this is the azeotrope conc. and it distills unchanged. . .

DetaDude - 18-12-2009 at 15:52

A "reversed Graham condenser" is not what Wilmad Labglass calls it.

And while we're at it why are we picking rat turds out of rice on a subject that I started this thread with. Read the first post as my intentions were to help some of our less learned citizen newbie chemists get a grasp on some of the various methods used to concentrate nitric acid, and share some of our experiments / experience. I did not wish to start a "I'm better than you" contest.

I just wanted to put some things in simple laymen terms,
..

hissingnoise - 18-12-2009 at 16:19

Anyway, DetaDude, you're lucky you can get 70% HNO3 to concentrate from H2SO4---most of us have to go the nitrate route for our strong acid, spending weekends washing KHSO4 out of flasks. . .
And I like the idea of not using up H2SO4 every time I want HNO3!
Still. . .

entropy51 - 18-12-2009 at 16:43

Quote: Originally posted by DetaDude

Has anyone tried concentrating nitric with a reflux set up,

Sooo, are you going to tell us what you mean by this? If you're using your relux condenser as a column, no you won't separate an azeotrope with a column.

[Edited on 19-12-2009 by entropy51]

hissingnoise - 19-12-2009 at 06:27

Gnnnge! What happened the Dulcolax reference?
It ain't something I'd need and even if I did, I'd 'pass' on it---prune juice for me if I needed anything.
But to the topic---to me, the HNO3 azeotrope is one of nature's cruelest cuts!
Why the fuck can't it be like ethanol?

User - 19-12-2009 at 06:36

Sure true!

Come on it really doesn't have to be that hard at all.
Why not buy 10 litre's of SA(battery) and a couple a kilo's KNO3
And a large flask (5 litre).
Take the weekend off !

Omg yes it is terrible

DetaDude - 19-12-2009 at 14:58

entropy51

Had you read my original post you would know by now that I am not distilling an azeotrope, which would be straight 70% HNO3 and that is why I mix H2SO4 with the nitric so when I run the distillation I'm dealing with pretty clean HNO3 vapor.

Does anyone but this guy understand what I'm saying , or can anyone else figure it out?

I believe hissingnoise has it figured out.

DJF90 - 19-12-2009 at 15:49

You're making very little sense DetaDude. If you're distilling something close to WFNA (as I suspect, under vacuum and using equal quantities of azeotropic nitric acid and conc. sulfuric acid) then whats the point in adding in a reflux condenser? To get anhydrous HNO3 you can distil the WFNA with another equal volume of conc. sulfuric... or something like that (I cannot remember the exact details).

DJF90

DetaDude - 19-12-2009 at 16:27

Thanks for the input. Just to clear the air I DO NOT use any reflux setup for the distillation of HNO3.

I do use such a setup when making booze, and have been told by others that they have used such a setup to distill HNO3 and since I've never heard of any results on this type of arraingment, I was hoping someone may of had some knowledge about this topic to share it with others.

Magpie - 19-12-2009 at 17:22

Deta Dude, does it strike you as being a little odd that all these experienced (hands-on) chemists can't figure out why or how you would use a reflux column for making nitric acid?

Why don't you provide us with a picture or a sketch.

[Edited on 20-12-2009 by Magpie]

magpie

DetaDude - 19-12-2009 at 18:57

What I find odd is that people on this and other threads start making comments without analyzing what the original/initial post was about.
Not all of the people on this board are graduate chemists, but merely average joe's/jane's with an interest in chemistry etc. so I for one would not be able to identify a "hands on chemist" if they bit me on the ass.

That said, on none of my posts have I made any claims to distilling nitric acid using a "reflux condenser", what I did say was that I've heard it reported elsewhere that this type of setup has been used. All I was asking is has anyone ever done this ,and if so how and what was the outcome.

As to a picture/diagram I'd be glad to do that , however I have absolutely no knowledge of how to perform the task of getting it from my desk into this computer and onto the discussion board.

I do use a packed column and reflux when I distill whiskey .

crazyboy - 19-12-2009 at 19:50

Quote: Originally posted by DetaDude

What I find odd is that people on this and other threads start making comments without analyzing what the original/initial post was about.
Not all of the people on this board are graduate chemists, but merely average joe's/jane's with an interest in chemistry etc. so I for one would not be able to identify a "hands on chemist" if they bit me on the ass.

That said, on none of my posts have I made any claims to distilling nitric acid using a "reflux condenser", what I did say was that I've heard it reported elsewhere that this type of setup has been used. All I was asking is has anyone ever done this ,and if so how and what was the outcome.

As to a picture/diagram I'd be glad to do that , however I have absolutely no knowledge of how to perform the task of getting it from my desk into this computer and onto the discussion board.

I do use a packed column and reflux when I distill whiskey .

Almost all of the regular members here are average people who experiment at home with chemistry, sure there are some that have made it a profesion but they also experiment at home or they wouldn't have much reason to be here at all.

As for none of your posts mentioning distillation with a reflux condesder, how about the one one the first page of this thread in which you say:

Quote:

Hmmmm reflux setup very close to what you just denied saying.

I suspect you mean fractional distillation which would require a column, this is possible however if you start with dry KNO3 and H2SO4 there won't be very much water in your product and any that is left could be eliminated by distilling your product mixed with additional sulfuric acid under a vacuum protected by a water trap made basic with NaOH.

A fractionating column isn't really helpful in this case because the only compounds present are HNO3, H2O and sulfuric acid (sulfuric acid won't distill easily.) Also nitric acid decomposes in the presence of heat and light which is why a fractionating column which returns vapor to the reaction flask would be counter productive.

Why are you so intent on getting nitric acid of such high concentrations anyway? nitric acid above concentrations of 90% is extremely dangerous and a nuisance to work with, I have some in my freezer in a Ziploc bag right now and even after frost forms on the outside of the bottle it still vents plumes of caustic vapors

Magpie - 19-12-2009 at 21:30

Quote: Originally posted by DetaDude

...what I did say was that I've heard it reported elsewhere that this type of setup has been used. All I was asking is has anyone ever done this ,and if so how and what was the outcome.

As to a picture/diagram I'd be glad to do that , however I have absolutely no knowledge of how to perform the task of getting it from my desk into this computer and onto the discussion board.

I do use a packed column and reflux when I distill whiskey .

>I tend to agree with crazyboy that those reports you heard of others using a reflux column were made by people who were using a fractionating column. Refluxing is inherent in a fractionating column's operation and maybe they were just not careful in their speech.

>Using computer technology by AOOF's can be a challenge! I should know. Do you have "paint" on your computer? (Start>All Programs>Accessories>Paint) If so, I recommend you learn to use it - it's fun. Get your kids or grandkids to show you how.

>For alcohol/water separation I'd say a fractionating column is essential. It's all got to do with relative volatility. Alcohol/water is a much tougher situation than water/azeotropic HNO3, especially if you start out with no extra water in the pot. You don't have that choice with fermented alcohol.

So I suspect your last sentence quoted above is inaccurate in the same way. Just say that you use a fractionating column. Mentioning reflux here just confuses the chemists.

Edit: I just thought of one other possibility: In industrial applications when using a fractionating column part of the condensate from the condenser is routed back to the fractionating column and added in as "reflux." This stream is usually added at the top of the fractionating column. It is used to obtain a condensed vapor of very high purity. Is this what you are doing with your whiskey?

[Edited on 20-12-2009 by Magpie]

[Edited on 20-12-2009 by Magpie]

hissingnoise - 20-12-2009 at 07:02

DetaDude said at the outset that he wasn't an experienced chemist, but now, what has transpired sounds like galloping, chemical 'political correctness'.
Why some people are making a big deal out of it is beyond me?
Chill out. . .

[Edited on 20-12-2009 by hissingnoise]

entropy51 - 20-12-2009 at 08:19

Quote:

Why some people are making a big deal out of it is beyond me?
Chill out. . .

As a charter member, you know that AOOOF are permitted a little grumpiness now and then...

hissingnoise - 20-12-2009 at 08:30

Ooof!. . .Touché!

crazy boy / magpie

DetaDude - 20-12-2009 at 08:40

Thanks for the input, your comments are well received. Not being a pro in the chemistry field or for that matter college level schooled in the subject, I will say I'm sorry for not making myself clearer, or using the proper or correct wording.

From the start of this thread I made it clear that I am not a pro grade chemist. All of my chemistry has been via the college of hard knowledge, a book here and there and hands on experiments.

Speaking of hands, I still have both of them and all ten of the original digits, given the fact that I was for some thirty years a journeyman machinist, and played around in the pyrotechnics field since age 11, at least I've been very cautious .

When I started this thread, I thought that I was pretty clear on my experience and that the thread was aimed at the novice people, and that I wished to keep it simple and "plain vanilla" I was hoping that the more astute chemists among the group would be able to understand what was being posted even if it seemed a little awkward or incorrect.

I'm sorry that I was somewhat short sighted about the ability of a very small percentage of those that post here, to use a bit of common sense.

I do want to say a sincere thank you for constructive contributions.

Hissingnoise.........Thank You

P.S. I have several grandkids that are computer geeks, but I live in a very rural area and they don't get out here very often.

User - 20-12-2009 at 09:17

Another approach that somehow seems very uncommon is the use af a dessicant.
For example P2O5 (which was proposed by someone i know from another forum) is a very suitable dehydrant.
He actually used it an proofed that it does work.
By adding enough to take away all the water in an azeotropic mixture one can obtain very close to 100% HNO3.
I know this is a very difficult chemical to obtain.
Funny how there are almost no references about this.

Another way i thought of is to use a XNO3 salt (preferably a heavier metal) that is clear of water.
By adding this to the nitric acid it could simply adsorb all remaining water.
Another issue here is that it isnt very easy to make these salt in anhydrous state due to their tendency to decompose instead of dehydrate.
An advantage of using XNO3 salts is that they don't react with nitric acid.
It could be quite straightforward and quick.

DJF90 - 20-12-2009 at 09:42

User: P2O5 is very well known for its use as a dessicant, and I would not hesitate to say it would most certainly increase the concentration of azeotropic nitric acid by reacting with the water present. However I disagree with "one can obtain very close to 100% HNO3". You CAN obtain 100%, and with even more P2O5, you can dehydrate the 100% nitric acid even further to its acid anhydride, dinitrogen pentoxide (N2O5), much like you can make SO3 from H2SO4 and P2O5.

Like you say, P2O5 is a very difficult chemical to obtain for most of us, so it makes sense to use conc sulfuric acid instead. Mixing azeotropic HNO3 with an equal volume of conc. sulfuric acid and distilling will get you to 90%+ HNO3, which can then be mixed again with an equal volume of conc. sulfuric acid and distilling to get essentially 100% HNO3. Then if you want to make N2O5, you waste much less P2O5 on absorbing water that can be removed from azeotropic HNO3 by non chemical means.

User - 20-12-2009 at 10:03

Sorry with uncommon i meant: using it to dehydrate HNO3
Or better: using P2O5 to dehydrate NA isn't mentioned very often.

DJF90 / User

DetaDude - 20-12-2009 at 10:15

I've heard of using P2O5 to pull the H2O from HNO3 but have never seen the process explained anywhere.

Would one of you or poss. another member be willing to write up a procedure for doing this, including ammounts, times, temps., etc.

Thanks in advance
,

gnitseretni - 20-12-2009 at 10:19

Quote: Originally posted by User

Another way i thought of is to use a XNO3 salt (preferably a heavier metal) that is clear of water.
By adding this to the nitric acid it could simply adsorb all remaining water.

Are you saying that, instead of distilling twice to get a higher concentration, I can just add some NH4NO3 to my say 90% HNO3, filter it, and end up with 95% or even better depending on how dry my AN is?

That sure sounds better than distilling twice!

crazyboy - 20-12-2009 at 10:36

Sorry, I don't see using phosphorus pentoxide to desiccate HNO3 as practical, from an economic standpoint or a chemistry one. As Wikipedia says:

"The desiccating power of P4O10 is strong enough to convert many mineral acids to their anhydrides. Examples: HNO3 is converted to N2O5; H2SO4 is converted to SO3; HClO4 is converted to Cl2O7."

So where have you heard of P4O10 being used to desiccate HNO3 DetaDude?

User - 20-12-2009 at 10:42

Quote:

[Quote] Originally posted by User
Another way i thought of is to use a XNO3 salt (preferably a heavier metal) that is clear of water.
By adding this to the nitric acid it could simply adsorb all remaining water.

Are you saying that, instead of distilling twice to get a higher concentration, I can just add some NH4NO3 to my say 90% HNO3, filter it, and end up with 95% or even better depending on how dry my AN is?

That sure sounds better than distilling twice!

No thats not really what i was saying.
If course it has to be a salt that 'captures' water, as far as i know AN does not have to ability to trap H20 in its structure.
I meant a substance that binds water to itself.

Does anyone know a method for preparing for example copper(||)nitrate in anhydrous form?
Simply heating it just decomposes the material.

[Edited on 20-12-2009 by User]

[Edited on 20-12-2009 by User]

DJF90 - 20-12-2009 at 10:44

If theres a limited quantity (calculated of course) of P2O5 then the water will react way before the HNO3. And any HNO3 molecules that do get dehydrated will react with the remaining water present to form HNO3 again, so theres no problem. Its only really economic if you want to make N2O5 as a nitrating agent, using it to dry azeotropic nitric acid is not.

You may be able to break the azeotrope by using a pressure swing distillation; I'm sure if you wanted to look information can be found in the context of distilling ethanol, perhaps check the Ethanol and Azeotrope pages on wikipedia.

gnitseretni: This wouldnt work. Like User says, you need the salt used to be hygroscopic, and generally anhydrous. Calcium sulfate may well work; although if you can find a nitrate salt that would work it would probably be a better choice.

[Edited on 20-12-2009 by DJF90]

gnitseretni - 20-12-2009 at 10:49

Quote: Originally posted by User

Ok, uhm.. well which one would you use? I'd like to give it a try the next time i distill me some HNO3, and if i can get whatever it is you suggest that is

hissingnoise - 20-12-2009 at 11:17

An anhydrous salt like magnesium nitrate can only be prepared from the anhydrous acid acting on the metal. . .
The same goes for copper!
The oxide produces a hydrate!
Anhydrous HNO3 can be prepared from dilute acid by dehydration by P2O5 using the correct stoichiometry!
The advantage H2SO4 has with dilute acid is that it's reusable after being reconcentrated.
Theoretically, distilling HNO3 from anhydrous Na or K nitrate and 98% H2SO4 should give 100% HNO3 but decomposition occurs and reduces concentration.

[Edited on 20-12-2009 by hissingnoise]

gnitseretni - 20-12-2009 at 11:25

Quote: Originally posted by hissingnoise

An anhydrous salt like magnesium nitrate can only be prepared from the anhydrous acid acting on the metal. . .
The same goes for copper!
The oxide produces a hydrate!
Anhydrous HNO3 can be prepared from dilute acid by dehydration by P2O5 using the correct stoichiometry!
The advantage H2SO4 has with dilute acid is that it's reusable after being reconcentrated.

I knew it sounded too simple to be true

User - 20-12-2009 at 11:38

I think Ca(NO3)2 could do the trick indeed.
It binds water
High melting point (561 degrees celsius) so I assume it can be baked to dryness..
It's hydrated form is the tetrahydrate so that means it can bind 4 water atoms per molecule.

Ca(NO3)2 =164.088 g/mol
Ca(NO3)2*4 H2O = 236.15 g/mol

1 gr Ca(NO3)2 can bind 0.4387 gr of H2O if my calculations are correct.

So if one has 10 gr of 70% HNO3 it contains 3 grams of water.
One needs to add 6.84 gr of Ca(NO3)2.

Dont know if it is practical but i think it can be done.

One could even use DCM to extract the HNO3 and distil off the dcm to obtain highly concentrated NA.
Just a mind spin.

[Edited on 20-12-2009 by User]

hissingnoise - 20-12-2009 at 11:48

Quote: Originally posted by User

Dont know if it is practical but i think it can be done.

Even here, distillation is needed as Ca(NO3)2 is soluble in HNO3.

crazyboy

DetaDude - 20-12-2009 at 13:36

Please show me where I stated that P4O10 was used to desiccate HNO3.

I think if you search the recent posts you'll find that I was refering to statements by others and nowhere did I mention P4O10.

Sorry.

entropy51 - 20-12-2009 at 14:45

Quote: Originally posted by DetaDude

I've heard of using P2O5 to pull the H2O from HNO3 but have never seen the process explained anywhere.

Quote: Originally posted by crazyboy

So where have you heard of P4O10 being used to desiccate HNO3 DetaDude?

[P2O5 is the empirical formula. P4O10 is the molecular formula. Same difference.

[Edited on 21-12-2009 by entropy51]

[Edited on 21-12-2009 by entropy51]

hissingnoise - 20-12-2009 at 17:09

It should be possible to distill ~98% HNO3 or better from H2SO4 consistently with an extra bit of care.
An air condenser between flask and water-cooled condenser might reconstitute some NO2 to HNO3, slightly upping concentration.
The receiver cooled below 0*C should help too.
But the main considerations are getting the nitrate salt completely anhydrous or as near it as possible and making sure H2SO4 is 98%.
Drying KNO3 by heat and popping it in a dessicator overnight would be good because what looks dry may still contain some little moisture.
Reconcentrating dilute H2SO4 is more difficult but done in dry/cold air with fume-extraction the finished conc. should be close to 98%.
Distillation then with a slight excess of H2SO4 should give near anhydrous HNO3.
Better than distilling twice. . . no?

crazyboy - 20-12-2009 at 18:01

Quote: Originally posted by DetaDude

Please show me where I stated that P4O10 was used to desiccate HNO3.

I think if you search the recent posts you'll find that I was refering to statements by others and nowhere did I mention P4O10.

Sorry.

http://www.fordhamprep.org/gcurran/sho/sho/lessons/lesson55....

per.y.ohlin - 20-12-2009 at 18:10

Quote: Originally posted by User

One could even use DCM to extract the HNO3 and distil off the dcm to obtain highly concentrated NA.
Just a mind spin.

[Edited on 20-12-2009 by User]

To avoid decomposition, the nitric acid can actually be precipitated out of the DCM. Carbonfeind describes the process here

hissingnoise - 21-12-2009 at 06:36

That procedure was first described on this board many, many moons ago!
Evaporating the DCM is probably easier as a simple ice-cooled receiver will condense DCM.
Welcome to SciMad, per.

User - 21-12-2009 at 07:28

It comes from a patent and has been discussed a million times on the web.
Pat. US 3981975
The whole point is that DCM isn't that good at dissolving NA.
One needs quite a lot of solvent to get some decent amounts of WFNA.

hissingnoise - 21-12-2009 at 07:36

Yeah, and the patent makes it sound a lot easier that the cold reality. . .
If you *can* get it out though, it'll be STRAAAWNG!

Ah the DCM method

DetaDude - 22-12-2009 at 08:56

I've tried this method, with only marginal success, it is somewhat time intensive, and yields are not to great. I was able to extract a small amount of WFNA >96% but it took awhile and plenty of DCM.

From an economic standpoint I'll stick with good ole distillation.

hissingnoise is correct about this method (DCM) it sounds a whole lot easier (from the patent) than it actually is in reality .

hissingnoise - 22-12-2009 at 09:06

Finding a seller willing to supply oleum would simplify things somewhat. . .

ScienceSquirrel - 22-12-2009 at 09:11

Getting your mitts on oleum is well nigh impossible.
Even if you can find a vendor who stocks it they want a lot more than it is really worth.

hissingnoise - 22-12-2009 at 09:19

We're back to oxidation by O3 circuitously. . .

hissingnoise - 22-12-2009 at 15:35

And to get a good yield from the DCM method the first essential is to shake the damn mixture until your two arms fall off. . .

ChrisWhewell - 22-12-2009 at 15:46

Any good chemist can make their own oleum when necessary.

entropy51 - 22-12-2009 at 16:00

Quote: Originally posted by ChrisWhewell

Any good chemist can make their own oleum when necessary.

Perhaps you'd like to share your favorite recipe. As you know, members here have done it, but it's not trivial.

hissingnoise - 22-12-2009 at 16:22

Yes, preferably something with 'Sulphur In'---> 'SO3 Out'---> would suit us best?

ScienceSquirrel - 22-12-2009 at 16:30

GPR sulphuric acid 98% ca £10 a litre.
40% oleum ca £75 a litre
And if you do not like the price you can always give it a go yourself.
Take a litre of sulphuric acid, build a sulphur trioxide generator and bubble the generated gas through the acid until the the solution has gained the requisite mass.
Easy enough

Ozone the ultimate oxidizer

DetaDude - 23-12-2009 at 09:16

Point of information:

Has anyone here actually used ozone for chemical synthesis ?

Say for example; to clean up HNO3, or to give SO2 encouragement to become SO3 et. al. etc.

If you have some experience with ozone please be so kind as to share your experiments with us.

hissingnoise - 23-12-2009 at 09:43

Quote: Originally posted by DetaDude

Has anyone here actually used ozone for chemical synthesis ?

I'd doubt it, DetaDude; its synthesis needs dry oxygen and its generation is woefully inefficient.
The cost of good equipment is very high, too!
A homemade generator can be cobbled together but it's unlikely to work continuously---or well!

hissingnoise

DetaDude - 24-12-2009 at 10:20

I was just refering to the use of ozone as a super oxidizer for use in other chem. syntheses.

I've overcome most of the pitfalls in the production of ozone . I just have not used it for chemistry stuff.

Many years back I developed an ozone generator, that runs on a 15KV neon transformer and now I'm looking to improve it with an oxygen concentrator, much like what they use for medical breathing therapy ..

hissingnoise - 24-12-2009 at 11:04

You're right DetaDude, I was being overly pessimistic. . .
Moist air in an ozone generator leads to the production of NOX and if the air-moisture could be controlled within tight parameters a mixture of NO2 and O3 might be formed.
After being given time to react the exhaust might consist of N2O5, theoretically, though some NO2/O3 would likely still be present.
All that's required then is dilute acid or water for absorption of the anhydride.
100% HNO3 will dissolve a sizable quantity of N2O5, leading to the nitric acid equivalent of oleum.
Of course it's easy to say. . .

DetaDude - 24-12-2009 at 17:01

I've found out through research on this subject that the time factor is more important than the moisture, but both play a part.

It seems that the speed at which the O2 (or air) moves through the ozone generator more or less governs which gas in made in the majority slow=nitrous fast favors ozone.

As soon as I get my hands on a concentrator and hook it up to the O3 unit I'm going to play around with some HNO3 and a few other things.

hissingnoise - 25-12-2009 at 05:02

A Jacob's Ladder could be configured in such a way that ozone would be produced along with NO2.
Output would be very small but running it continuously would build product over time.
I fear though, we may be clutching at straws.
KNO3/H2SO4 is triedand trusted.

malford - 1-10-2013 at 21:26

Some searching brought this thread to my attention, though the first couple of pages brought me disgust. It's opening post is exactly the discussion I am looking for. However, I must preface my post with the declaration that crazyboy, entropy51, bbartlog and others involved are the top three least sensible, to put it kindly, individuals whose writings I have read or ever will read. The unlikely alternative, of course, is that he or she each are exceptionally skilled at trolling.

Delta specifically stated he was looking for anecdotes from using a reflux condenser (in other words, a condenser that is designed to be used for reflux) for distilling nitric acid with the rest of the setup being that of a typical distillation. This is something I have wondered about myself. In the unlikely event that he meant using an intact reflux setup for distilling nitric acid, even then, an ounce of brain matter should be sufficient to conceptualize how this would work. I have even thought of using this myself previously. Examine this part. The boiling flask would be connected to the lower, right inner joint. The reflux condenser would be connected to the upper, left outer joint. A collected flask would be connected to the lower left stopcock.

More on topic, 70% nitric acid and 98% sulfuric acid are very readily and cheaply available to anyone in the US and probably elsewhere. As such, I am distilling the nitric acid from a mixture with the sulfuric acid. This seems to be one of the most effective approaches available to me for highly concentrated nitric acid. I am looking for more information on what exact compounds are formed when the 30% water from the nitric acid is reacted with the 98% sulfuric acid. Will this newly formed compound be azeotropic with nitric acid? If so, at what point is the azeotrope?

testimento - 7-10-2013 at 16:02

One shall do following reaction:

1) Roast some calcium carbonate to calcium oxide
2) Mix it with water to get hydroxide
3) Get some otc fertilizer (active ingredients are mostly AN and urea), dissolve'n'filter'n'evaporate
4) Mix CaOH into it at saturated, near-boiling water solution and lead evolving ammonia into water to recover it
5) Boil off the residual water and heat up to 500C and take the evolving ammonia and isocyanic acid fumes by leading them into water trap
6) Roast the residue, which is calcium nitrate, and it will decompose into nitrogen dioxide and calcium oxide
7) Lead NO2 fumes into water to get nitric acid up to 68%
8) Perform an azeotropic vacuum distillation to acquire WFNA
9) Reuse the CaO

1kg of calcium carbonate with 1600g AN will produce 1260g of 100% HNO3 at 100% efficiency, practically yields of 800-1000g should be achievable.

For lazy bastards, mixing hydrochloric acid with ammonium nitrate will yield ammonium chloride and nitric acid.

Oscilllator - 7-10-2013 at 22:42

Quote: Originally posted by testimento

7) Lead NO2 fumes into water to get nitric acid up to 68%
8) Perform an azeotropic vacuum distillation to acquire WFNA

What??
Firstly, you can theoretically achieve 98% nitric acid by the absorption of NO2 vapors. This is how its done on an industrial scale, although admittedly it probably isn't practical for the home chemist. (source: wikipedia)
Secondly, I fail to see how an azeotropic vacuum distillation will get you above the azeotrope... assuming that by WFNA you mean nitric acid of a concentration greater than 95%? (definition of WFNA taken from wikipedia)
Also, do you have any sources for the production of calcium nitrate using that method?

DetaDude - 8-10-2013 at 15:28

Malford..................Thank you for your kind understanding, and your ability to understand the Kings English.

Some of those that respond here do not seem to have a very good grasp of common sense and understanding. Hissingnoise has a great deal of comprehintion and you also seem to "get it".

Thanks again DetaDude .

malford - 8-10-2013 at 16:24

I can now answer my own question from my previous post above.

Examine this: http://www.akersolutions.com/Documents/PandC/Mining%20and%20...

On page 2, the chart shows that 50% by weight of sulfuric acid added to the nitric acid eliminates the azeotrope. The question then becomes, what is the vapor composition with this mixture? Using the answer to that question, we can determine whether a single distillation will suffice or many theoretical plates are needed.

If anyone has or knows how to figure what the vapor composition is of a ternary mixture of nitric acid, water and sulfuric acid, then my appreciation would know no bounds.

Magpie - 8-10-2013 at 19:03

The sulfuric acid is basically non-volatile at the temperatures used to distill over the nitric acid. Measure the density of your product and by using a handbook table you will find the concentration. Or titrate it.

testimento - 9-10-2013 at 22:16

Quote: Originally posted by Oscilllator

Firstly, you can theoretically achieve 98% nitric acid by the absorption of NO2 vapors.

Secondly, I fail to see how an azeotropic vacuum distillation will get you above the azeotrope...

Also, do you have any sources for the production of calcium nitrate using that method?

Alkali mixed with ammonium nitrate will yield ammonia and calcium nitrate.

Azeotropic distillation is performed mixing the nitric acid with sulfuric acid, as stated. The SA is recoverable by boiling off the water.

I noted the practical yield being 60-80% max.

Anyway, that's how one can get either dilute HNO3 up to 68% without SA, or conc. HNO3 with SA, both ways without wasting any acid.

Many nitrations can be carried out using fuming nitric acid, so I don't see any single practical reason to waste - for most people - precious sufuric acid for it, if it can be avoided. For ex. PETN synthesis yields quite an equivalent amounts of product with, and without.