Sciencemadness Discussion Board

concentrated nitric acid from 'strong' acid and H2SO4

Antwain - 11-9-2007 at 06:39

Ok, first of all, I know that it may be better to turn my nitric acid into the sodium salt, stick that in the oven and then distill with conc. H2SO4, but I really hate to waste chemicals so I will explore this first.

An old inorganic book I have, can't remember which and cant find it just now, tells me that 98% nitric acid can be made from 68% nitric acid and 98% sulfuric acid. However it didn't give any further info on this one. Ive seen it somewhere else too (probably wiki, so its not worth anything and casts doubt on the whole thing :D )

Does anyone have any further info on this procedure, such as the relative amounts of the two acids used, does it need to be refluxed at low temp (ie just above the boiling point of anhydrous nitric) or whether it gives good yields or is only good for industrial scale manufacture (I think this was where it was supposedly used).

Alternatively can anyone suggest any good methods for making conc. nitirc acid.

One other thing.. anyone who has owned/used conc./anhyd nitric acid - how quickly and or readily does it decompose into N oxides? I understand that it is somewhat stable if kept away from light?

Sauron - 11-9-2007 at 07:59

It depends on the concentration of your nitric acid, as indicated by its density, and to measure that you either need a range of hydrometers from about 1.0 to 1.6 (g/cc)

Or you need a good volumetric flask and a scale that will allow you to tare the flask and weigh it when full to the mark. 500 ml or 1 L will do. Smaller than that and your precision gets fuzzy.

Once you know how strong your acid is to start with you can determine how much conc H2SO4 is needed to remove the water.

This procedure is normally used to prepare very concentrated (fuming) nitric acid up to 100%. See Brauer's book in forum library. Naturally you can then dilute your acid back to ordinary concentrated d 1.42 if that is your goal.

garage chemist - 11-9-2007 at 08:06

The procedure is in Brauer, and in short form on orgsyn:
Quote:
The anhydrous nitric acid (sp. gr. > 1.53) is most easily obtained by slowly distilling ordinary fuming nitric acid from its own bulk of concentrated sulfuric acid. If ordinary concentrated nitric acid (sp. gr. 1.42) is used, it is advisable to distil twice from equal volumes of sulfuric acid.


From http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv3...

I have made fuming HNO3 from 68% and even 53% for as long as I can remember, you need to distill it twice from its own bulk of conc H2SO4. NO refluxing, and NO column! Both would decompose your HNO3 much more than necessary.

In the first distillation, collect everything until the temperature reaches 121°C, and in the second one, ONLY collect the fraction coming over at 84-88°C (discard forerun, this is mainly nitric oxides). This is your 99% HNO3. It will be yellow of course, due to slight decomposition during distillation as is always the case when fuming HNO3 is distilled at atmospheric pressure.

For decolorizing it (if really necessary), first redistill in vacuum from a small amount of H2SO4, then do the well- known treatment with urea or air bubbling. The acid wont stay colorless for long. The decolorizing also decreases the concentration a bit.

Also, concentrated nitric acid is 68%! The 99% acid is called fuming nitric acid, not concentrated.

Antwain - 11-9-2007 at 08:17

Quote:
Originally posted by garage chemist
Also, concentrated nitric acid is 68%! The 99% acid is called fuming nitric acid, not concentrated.


Are you sure about that? That is what I also thought (at least that 68% was 'concentrated') but I have now seen two reactions that called for "concentrated nitric acid (specific gravitity 1.5)" and "concentrated nitric acid 98%" also F.Sherwood Taylor says "These include fuming nitric acid, which contains dissolved nitrogen peroxide and is yellow or red in colour, and 'concentrated nitric acid' of density 1.5 containing about 98% HNO3. The ordinare 'strong nitric acid' is of density 1.4 and contains about 65% of HNO3."

And thanks for the info.

Sauron - 11-9-2007 at 09:30

The standard reagent grade is 68% and d 1.42.

Nitric acid of d. 1.5-1.52 is fuming nitric acid, either red fuming (containing dissolved NOx) or white fuming (which has had the dissolved oxides removed or destroyed chemically. Nitric acid of such concentration is VERY expensive and is also costly to ship due to hazmat regs and charges. It is a very powerful oxidizer, and in fact is a component of hypergolic liquid rocket propellants.

So it is fairly common to prepare it as needed from the cheap(er) reagent grade.

If you are audacious enough to question garage chemist, I suggest you look in Merck Index, or the ACS reagent specification for HNO3, or both. Your references must be ancient. I've been a chemist for more than 40 years and concentrated nitric acid has been 68% and d 1.42 for a lot longer than that.

[Edited on 12-9-2007 by Sauron]

garage chemist - 11-9-2007 at 10:31

About the price thing: You can prepare about 1L 99% HNO3 from 2L 68% HNO3 and the sulfuric acid is completely reusable (at least for this purpose) if you reconcentrate it (the residual HNO3 boils off before the water).
The 99% HNO3 is really best made yourself from 68% acid instead of purchased, due to the huge money savings and the limited storage stability.

And since the H2SO4 is not used up it is even better than the NaNO3 method.

The term "concentrated nitric acid" does not seem to be well defined in english literature, with different authors meaning different things as Antwain showed with his citation, but in german literature it sure is, concentrated always meaning 68% with HNO3.

I remember that there once was a misunderstanding here with HNO3 concentration, with someone having used fuming HNO3 instead of concentrated and the synthesis producing something different than desired:
http://sciencemadness.org/talk/viewthread.php?tid=4240&p...

99% HNO3 is completely different in its reactivity than 68%, with 68% being compatible with plastics like PP (bottle screw caps) and 99% not being compatible with any plastics except fluorinated ones like teflon. A bottle or cap that holds up perfectly to 68% HNO3 may well be destroyed by 99% HNO3 in short time. Use only glass bottles and teflon lined caps or ground glass stoppers.

Antwain - 11-9-2007 at 13:34

Double thanks for the warning about plastics. I was going to make it today and put it in a glass bottle with a teflon thread but a not-teflon lid. Now I know to transfer my "crud nitric acid" - pure, but percentage unknown and not too high to that bottle and put the fuming nitric into the glass stoppered quickfit bottle. That also probably explains why it took 30 long, painful hours to get a crap yield of some iodic acid - the book specified conc. but went on to say sp. grav 1.5, which I missed... It probably would have gone much faster with 98%.

Also, does that mean that i *shouldn't* use 98% acid with sulfuric for a nitration(ester)? I was going to do that at some point because I figured that since it apparently dosent work too well once the water content gets up to 30% that I could save some reagents if I used fuming nitric acid... maybe not then :(

Sauron - 11-9-2007 at 15:11

What book(s) are you using?

And which ester are we talking about?

In general if the book called for conc nitric acid I would not substitute fuming. Was density (sp gr) specified?

Fleaker - 11-9-2007 at 16:15

Be careful with fuming nitric acid, it will explode very easily with some organics (like ethanol). Garage chemist is very correct in saying that it will quickly and reliably destroy phenolic caps, and most other plastics. If a reaction calls for concentrated nitric, don't assume that it is fuming. If it didn't specify density, I would use the 68-71% stuff.


I store several liters in this type of bottle:
(hijacked from an older post) http://img231.imageshack.us/img231/6012/hno3inteflonal8.jpg

ordenblitz - 11-9-2007 at 16:21

I got lucky enough to find two of these same solid PTFE bottles on EBAY.
Very expensive even second hand and nice to hold in your hands, heavy and cold.
;)

Fleaker - 11-9-2007 at 16:31

Hahahah, I see you are smitten with them as well. They are absolutely friggen amazing. I have I think 15 or 16 of those 2L bottles and some PFA and PTFE beakers too. If I could find them cheaper, I'd buy many more. They hold up so well to heat, pressure, and chemicals!

For storing HF, xSO3.H2SO4, fuming nitric, and hydrazine, they can't be beat!!

not_important - 11-9-2007 at 17:40

garage chemist is correct that the term "concentrated nitric acid" has had variable meaning in English language books. I've seen changes in the meaning as you move from the mid-Victorian era on up to a bit pre-WWII, also between British and US writings, and in works targeted at different industries - the explosives and the dye industries in particular.

So if you're referencing an older work, you may need to do some extra work to ascertain the meaning of "concentrated" for that particular work.

chief - 14-9-2007 at 06:34

The not-so-complicated way: Ba(NO3)2 + H2SO4 ==> HNO3 + BaSO4, where the BaSO4 is to be found as insoluble solid at the ground. At least that's what I theoretically know about it ...
I might like to know how it actually works.

[Edited on 14-9-2007 by chief]

Antwain - 14-9-2007 at 07:23

Without a doubt the best way to do it.... If you have kilos of Ba(NO3)2 you want to turn into useless BaSO4, and some glass wool or a sinter to filter it through. I, however, have several 10s of grams of Ba(NO3)2 and only a few cc of glass wool, neither of which I particularly want to frivolously use :D

I think that the Ba salt + H2SO4 method is best saved for making HMnO4 of HClO3 not that I have the required salts :( Deffinately things on my 'to do one day' list, though.

chief - 14-9-2007 at 12:20

That HMnO4- synthesis is very interesting for me; I have no HClO3, but some NaCl/NaClO3 solution from electrolysis. I would try Ba(MnO4)2 + H2SO4 ==> 2 HMnO4 + BaSO4 ;
the Ba(MnO4)2 I have stems from Ba(NO3)2 + KMnO4 ==crystallizing==> Ba(MnO4)2 + KNO3 ;
but I still search for a better way (without the heating/cooling/watching etc. of the recrystallization)

chief - 14-9-2007 at 12:23

I forgot above: I just want the Ba(MnO4)2 ; therefore the HMnO4 from HClO3 would be interesting for me (then 2 HMnO4 + BaCO3 ==> Ba(MnO4)2 + CO2 + H2O), if easier than what I wrote above ...

Antwain - 15-9-2007 at 04:08

What I meant was that If you have Ba(ClO3)2 then you can make HClO3 in a cold, dilute solution. It is not stable and can only be made, in the cold, and used straight after. Presumably you could make the Ba(ClO3)2 from electrolysis of BaCl2 solution, or buy it from somewhere. I have no idea how one would go about making Ba(MnO4)2... I only know that a dilute solution of this can be used to prepare a dilute solution of HMnO4, which in turn can be used to make other permanganates.

chief - 16-9-2007 at 06:33

Ba(MnO4) can be made as I said above (although I search for a better way), and no HMnO4 needs to be prepared, just the
(Ba(MnO4)-solution) + (sulfate-solution) ==> BaSO4 + (permanganate of the other ion (from the sulfate))
That I know theoretically, but it should really work.