A solution of HNO3 and NH4Cl will not nitrate organic compounds such as starch, sugar, glycerin, etcetras.plasma - 20-5-2002 at 18:50
Damn!
Thanks anywayplasma - 16-6-2002 at 12:08
I think I saw a post here (this forum) somewhere which had a list of acids which could be used as a subsitute for H2SO4 when nitrating. Could someone
give me the link (or post it) because I can't find it.
I am asking this because the only way to get H2SO4 is to buy it as car batteries and destill. The problem is that car batteries are rather expensive.
(I don't have very much money)
I can get H3PO4 (30 %) as a rust remover though, maybe that will work.
Thanks ! Rhadon - 17-6-2002 at 11:36
But you don't have to buy the whole batteries! Just go to a garage and ask if you can have some battery acid because you need it for experimental
purposes.
If you don't want to tell them that you'll use it for chemical experiments, you could say that you wanted to try building a battery with it.
I never had problems with obtaining sulfuric acid this way, and I never had to pay for it, too.
[Edited on 17-6-2002 by Rhadon]plasma - 17-6-2002 at 12:06
Thanks Rhadon I might just try that.
BTW, does anyone know if H3PO4 can be used instead of H2SO4 to nitrate ?madscientist - 17-6-2002 at 12:47
I highly doubt it will work.vulture - 20-6-2002 at 11:13
I have tried to nitrate cellulose with 85% H3PO4 and 70% HNO3, but it failed. Of course cellulosehexanitrate is a very unforgiving compound when it
comes to nitrating.
You can't make HNO3 from H3PO4 as it is a much weaker acid, but you could try it while heating the reagents.plasma - 24-6-2002 at 14:39
I finaly found the post I was talking about, it was posted by Coen.
Quote:
Other acid-catalyzed nitrations can be done using mixtures of HNO3 with HF, HClO4, BF3, TFA, TFAA etc.
Obviously this cannot be used to nitrate cellulose because the H2SO4 is needed to "dehydrate" it. Does anybody know what compounds any of theese mixes
will nitrate (HClO4 and HF in particular)vulture - 25-6-2002 at 06:38
I wouldn't nitrate any flammable substances with a HClO4/HNO3 mix! The biggest risk is spontaneous ignition and secondly, some highly unstable
perchlorate esters will form. plasma - 25-6-2002 at 08:50
Thanks Vulture !
Any thoughts about the HNO3/HF mix. (ratios to use, dangers etc.)madscientist - 25-6-2002 at 09:34
Hydrogen fluoride is extremely toxic.vulture - 8-7-2002 at 01:13
I find that HF as nitrating catalyst highly unlikely, except a few resistant organic molecules, everything else will fluorinated in stead of nitrated.AndersHoveland - 20-9-2011 at 13:29
Solid NH4NO3 dissolved in 30% concentrated HCl is likely not entirely stable over time. There are various complex equilibria going on, and the
ammonium ions potentially could slowly be attacked.
(3)NH4NO3 + (6)HCl --> (9)H2O + (3)N2 + (3)Cl2
Here is various information about different equilibria and reactions to substantiate this idea.
It is known that there exists the following equilibrium in aqua regia:
HNO3 + (3)HCl <==> NOCl + (2)H2O + Cl2
Now, you might argue that this only happens because 70% nitric acid is being used, but even 30%HCl with NaNO3 gives off brown fumes when heated, and
is capable of attacking gold.
It is also known that a mix of nitric oxide and nitrogen dioxide, together can oxidize ammonium cations. This apparently has more to do with an
ill-understood reaction mechanism, rather than oxidizing strength. A mixture of nitric oxide (NO) and nitrogen dioxide (NO2), obtained from the action
of sulfuric acid on sodium nitrite, was bubbled into a boiling solution of ammonium perchlorate. Although the reaction was very slow, some of the
ammonium ions were oxidized, leaving free perchloric acid in solution.
Chlorine gas may be bubbled into an aqueous solution of ammonium perchlorate for several hours with almost no observable reaction. Actually the
reaction between Cl2 and ammonium salts is an equilibrium reaction.
NCl3 + (4)HCl <==> NH4Cl + (3)Cl2
The formation of oily drops of NCl3 is only favorable within a narrow pH range (around pH6), however.
Neither does concentrated nitric acid oxidize ammonium perchlorate.
It was observed that pure nitrogen dioxide oxidized ammonium perchlorate such that free perchloric acid could be obtained. The addition of nitric
oxide, or if oxygen, to the nitrogen dioxide did not observably efect the reaction. Indeed, no reaction was observed when pure nitric oxide was
bubbled into ammonium perchlorate. It was also observed that more than two molecules of NO2 formed for every molecule of perchloric acid. A.W. Knight. Undergraduate Thesis at California Institute of Technology (1992). "Perchloric Acid from Ammonium Perchlorate and Oxides of
Nitrogen
“reaction of ammonium perchlorate with nitric and hydrochloric acids, and then concentration at 198–200C to eliminate the unreacted acids by
vacuum distillation:
Ironically, slowly and gradually adding a small quantity of a reducing agent can allow concentrated nitric acid to oxidize ammonium ions. If formic
acid is slowly added into a boiling solution of ammonium perchlorate and nitric acid, all of the ammonium ions in solution can be completely
oxidized. Both of these reactions are thought to be due to the intermediate formation of nitrous acid, HNO2, which is capable of oxidizing ammonium
ions. If, however, hydroiodic acid is slowly added to the boiling solution of ammonium perchlorate and nitric acid, the ammonium perchlorate is not
decomposed to any significant extent. It appears that although the hydroiodic acid initially reduces the nitric acid to form nitrous acid as an
intermediate, the nitrous acid is reduced by the hydroiodic acid at a much faster rate than it is formed. The oxidation of ammonium salts using nitric
acid proceeds much more rapidly, and gives higher yields, when hydrochloric acid is used, rather than formic acid, although more concentrated nitric
acid is required when hydrochloric acid is employed.
I also had the idea of "forcing" the reaction between ammonium perchlorate and chlorine, by subjecting the reaction to ultrasonic agitation (such as
from a common mystifier). This might cause decomposition of the small portion of NCl3, that exists in equilibrium, into N2 and Cl2, potentially
shifting the reaction forward.
