I think molten AN could be electrolysed to give a mixture of nitrogen oxides and oxygen at the positive electrode (more than enough oxygen to give
HNO3 by dissolving in water) and ammonia + hydrogen on the negative electrode.
Could someone try it and also suggest a suitable anode (to resist hot nitrate and hot-nitrogen-oxides corrosion) apart from platinum?vulture - 24-6-2003 at 08:58
Ammoniumnitrate decomposes very quickly when molten. Secondly, would molten ammoniumnitrate conduct electricity? It's not really an ionic salt.Polverone - 24-6-2003 at 12:54
Ammonium nitrate doesn't decompose very fast if you hold it just above its melting point. Still, if you want to make nitric acid, I think it
would be a lot more practical to convert it to sodium nitrate (heat with NaOH or Na2CO3) and distill that with sulfuric acid.
AN-»nitric acid + ammonia
Theoretic - 25-6-2003 at 03:55
"It's not really an ionic salt."
What do you mean?
What, apart from an ionic compound with the formula of NH4+ NO3- can AN be?vulture - 25-6-2003 at 08:37
You don't have to get angry at me.
What I was trying to say is that ammoniumnitrate is not an ionic salt like NaCl, because of the N-O-N bond, which has very little difference in
electronegativity. Therefore it could be possible that it does not conduct electricity when molten.
That's just a suggestion I was making and I don't see why it was necessary to flame me for that.Theoretic - 26-6-2003 at 06:52
"...because of the N-O-N bond, which has very little difference in electronegativity..."
Even if atoms making up the nitrate ion(/group) have little difference in electronegativity, that doesn't make it AN nonionic.
The anion doesn't need a dipole moment to make the compound ionic.
Maybe you mean that NH4+ makes a covalent bond with NO3-? It isn't possible, because there are only single bonds in NH4+, so to make a covalent
bond it would have to break up.
The structure of AN:
O O - H +
\ / |
N H--N--H
|| |
O HTheoretic - 26-6-2003 at 06:54
Sorry about the drawing, it got crunched on its way by the computer.Theoretic - 30-6-2003 at 06:58
Why can't you distill AN with H2SO4 to get HNO3 and (NH)2SO4?rikkitikkitavi - 30-6-2003 at 09:07
if AN wasnt a ionic compound , aqeous solutions of AN wouldnt conduct electricity as well as it does. I m positively sure that molten AN conducts
electricity, I havent heard of any compound showing conductivy in solvatized state (dissolved in water f e x ) and not conducting electricity in
molten state.
And yes, HNO3 can be made by distilling AN+ H2SO4. The process is however more common with Na,K and hence more described in detail.
/rickardTheoretic - 1-7-2003 at 06:40
Sorry about the question on AN + H2SO4 distillation, I should have realised that (NH4)2SO4 could be soluble in conc. H2SO4 and contaminate the
product.
And would someone suggest a NOX-and-hot-NO3- resistant anode for the process? rikkitikkitavi - 3-7-2003 at 10:13
since you distill of HNO3 , it will not be contaminated by (NH4)2SO4, but the latter salt starts to decompose at 100 C, into NH3 and NH4HSO4,
decomposition increasing heavily above 120 C, which could be one way of explaining lower yields of HNO3 compared to Na;K .
suggestion for anode: our all time favourite : Pt!!
/rickard
[Edited on 3-7-2003 by rikkitikkitavi]CROpyrO - 1-8-2003 at 08:40
if you use iron electrodes will it during the electrolyse form iron(II)nitrate-hexahidrate ?chemoleo - 4-8-2003 at 03:56
probably not unless you do the electrolysis in an aqueous solution. If you electrolyse with molten AN, you will probably get gases. but N2O will be
produced with molten AN anyhow, electrolyis or not.Theoretic - 4-8-2003 at 06:56
Er, chemoleo, there's actually a 40C gap between AN's melting temperature (170C)
and its decomposition point (210C).blip - 4-8-2003 at 18:53
The temperature distribution curve for ammonium nitrate at 170°C likely has at least some of the liquid at the decomposition temperature and
decomposing. Water doesn't just evaporate at 100°C+, instead it does so gradually more and more as the temperature is raised.
[Edited on 8-5-2003 by blip]kryss - 6-8-2003 at 15:29
1. Heating a nitrate with an ammonium salt ( even in solution ) oxidizes the ammonium ion so you wont get nitric acid distilling off from sulphuric
& A.N - Nitrous oxide is one of the products. IF you add some alcohol to this mixture you'll get ethyl nitrate or nitrite coming off. We do
this to get rid of the nitrate if we're analysing for ammonia - but in small quantities & not in the presence of lots of nitrate.
2. AN is ionic - electrolyse it with a mercury cathode and the NH4+ gains an electron & dissolves in the Hg - Ammoonium amalgam cf Na & K.
chemoleo - 6-8-2003 at 17:08
kryss, are you saying this is a route to Et-NO3 synthesis (or methylnitrate?)? how exactly can this be done??kryss - 7-8-2003 at 15:43
Think you'd have to mess about with concentrations to get an appreciable yield - the levels we would be working at you'd just get a yellow
colouration in the sample - and doing it at high concetrations is dangerous - alcohol and nitric acid tend to be explosive combinations although in
fairness the context in which i was reading about this was to do with mirror silvering & you can get silver fulminate forming.Theoretic - 8-8-2003 at 05:49
Actually, ammonium amalgam decomposes readily at ordinary temperatures. It COULD be that you still get ammonia and hydrogen, but don't hope to
make HNO3 that way: ammonia is absorbed instantly by water (the so-called "ammonia fountain", never mind acid.axehandle - 19-2-2004 at 16:40
Quote:
And would someone suggest a NOX-and-hot-NO3- resistant anode for the process?
Platinum or iridium should not be attacked unless you have chlorine in there somewhere and are using AC.
Edit: I suppose gold would work as well, and be considerably less expensive as well.
/A
[Edited on 2004-2-20 by axehandle]
KNO3 Electrolysis Attempt
hodges - 23-2-2004 at 16:37
Here is what happened with my (not very successful) attempt to produce HNO3 (and KOH) through the electrolysis of an aqueous solution of KNO3.
I dissolved 26 grams of KNO3 in each of two cups containing around 160ml of water. I used a moist paper towel to join the cups, and performed
electrolysis with carbon rods from an old lantern dry cell. I calculated that with 7AH of current, there should be 0.5 moles of HNO3 produced as well
as 0.5 moles of KOH, giving 3M solutions of each. I ran the electrolysis for 24 hours. Current ranged from 220mA at the start up to 360mA at the
end.
I noticed from the beginning that more bubbles were being produced on the positive (oxygen) side than on the negative side. This was unexpected,
because twice as much hydrogen as oxygen should be produced. As the reaction proceeded, the rates of gas production on the negative side compared to
the positive side increased, but was still slightly less. Early on there was a nitric acid smell on the positive side. As the reaction proceeded,
though, this smell actually seemed to decrease. By the end of the reaction, there was a noticable ammonia smell on the negative side.
When I stopped the electrolysis after 24 hours, the positive carbon rod had mostly been eaten away. The negative rod had not eroded at all. I
titrated both solutions with a known acid/base and found that both were around 0.8N, compared to the theoretical ideal of 3N.Saerynide - 27-2-2004 at 07:53
Im confused... If NO3- is an anion, why would it end up at the cathode as NH3? How does it get reduced? Wouldnt it be attracted to the anode? hodges - 27-2-2004 at 15:53
Quote:
Originally posted by Saerynide
Im confused... If NO3- is an anion, why would it end up at the cathode as NH3? How does it get reduced? Wouldnt it be attracted to the anode?
There's NO3- on both sides - that's the idea of electrolysis (to separate it). The cathode would be the side where reduction occurs. Also,
since that side is basic, it would react with any ammonium salt produced to make NH3. Although it was not exceptionally strong, there is no question
in my mind that I smelled NH3, and it was definitely on the negative side.Saerynide - 27-2-2004 at 20:03
Oh right, I forgot it was on both sides.. hahaha silly me.Quince - 10-2-2005 at 02:16
As someone mentioned distilling nitric acid from ammonium nitrate and sulfuric acid previously in this thread, I feel I can ask about it too.
I've read that it is not as good a method as when using sodium or potassium nitrates due to decomposition of the ammonium nitrate. My question
is, can't this be avoided with vacuum distillation?neutrino - 10-2-2005 at 03:27
Yes, that's the standard method.gtacchi - 22-2-2005 at 04:55
The better way to produce HNO3, is mix carefully sulphuric acid in NaNO3 with heat and cool the fumes, you produce a nitric acid, with a minimum
contamination.neutrino - 22-2-2005 at 14:19
Yes, that’s the standard <i>cheap</i> method, it has been discussed extensively, especially at RS.
, never mind acid.