I was reading, and saw that ammonia + lithium => lithium amide. And that burning the amide releases nitrogen dioxide.
Is this what actually happens? Seems like it might be a way to produce nitric directly from air without having to use corona discharge.
And that Lithium when powdered and reacted to dry air will form Li3N (grabbing the nitrogen out of the air), which produces ammonia when exposed to
water.
As well Lithium nitrate decomposes to Lithium oxide and NO2 and O2.
So putting this all together it seems lithium would be a good catalyst for making nitric acid from nitrogen in the air.
a) it can produce its own NH3 from the atmosphere, via Li3N + 3H2O -> NH3 + 3LiOH
b) the 2Li + 2NH3 + -> 2LiNH2 + H2
c) burning 2LiNH2 + 4O2 -> Li2O2 + 2NO2 + 2H2O
and even if there ends up being LiNO3 produced it is decomposed back to NO2:
d) 4LiNO3 -> 2Li2O + 4NO2 + O2
I've searched the forum but found nothing on this.
If all of this is true, there will probably be a shortcut that can be done, or at least a clever setup to make nitric acid using lithium.
Am I way off base here? Should I start thinking up some experiments to do?watson.fawkes - 20-10-2008 at 20:14
I don't see a closed lithium cycle. You've got unbalanced Li metal on the left and both LiOH and Li<sub>2</sub>O on the right. Or perhaps
you don't actually mean "catalyst", even in a broad sense.
Even if I grant that you can recover the metal by electrolysis, you'd need to make an argument that (1) this is better on either (a) fixed equipment
costs or (b) variable energy costs and (2) that there's some regime where the advantages of one outweigh the disadvantages of the other. I think (1a)
is likely false. Point (1b) might be true, that the energy efficiency of electrolysis is better than corona, since heating losses might be less, but
this isn't immediate for me. And even if it's true that (1b) outweighs (1a), that puts the applicable regime in the industrial production category,
where the minimum production volume (<i>per annum</i>, say) to hit the tradeoff point is likely to be fairly large.indigofuzzy - 21-10-2008 at 10:53
Maybe the method isn't more efficient, per se, but more accessible to OMG.
It may well be worth exploring, as Nitric acid is a useful reagent, and we may as well have as many useful syntheses as possible.watson.fawkes - 21-10-2008 at 11:33
Quote:
Originally posted by indigofuzzy
Maybe the method isn't more efficient, per se, but more accessible to OMG.
It may well be worth exploring, as Nitric acid is a useful reagent, and we may as well have as many useful syntheses as possible.
I am not advocating suppression of this technique or its exploration. I am advocating ensuring that it's
actually better under some measure, any measure at all. Corona discharge is pretty easy to set up. To my eye it's easier than electrolysis of lithium
compounds. Make the argument; that's what I'm asking for.chemoleo - 21-10-2008 at 15:03
OMG, please explain why you think that Li metal is an economic resource, industrially or in home labs. From batteries???
THen, the reaction to the amide needs to be done with anhydrous oxygen-free NH3, possibly at higher temperatures, and NH3 needs to be recycled to
obtain any sort economical reaction.
Then, LiNO3 decomposes only at >800 deg C.
The resultant Li2O or peroxide needs to be converted back to the hydrochloride (or whatever is best for electrolysis), and electrolysed (whilst the
chlorine gas is another issue).
A huge expenditure of energy and effort - there are FAR easier ways to get NO2 gas!!!!kclo4 - 21-10-2008 at 16:29
If you have NH3 to begin with, Why not just try to oxidize that to HNO3 with some sort of Nickel/Copper catalyst? That is an exothermic reaction and
those metals are easy to acquire.OMG - 21-10-2008 at 16:32
First, I am by no means a chemist looking for a more economically industrial route to nitric acid. I am simply interested in chemistry (and learning
new things every week).
I was reading about lithium (previously I was learning about nitric acid synthesis) and it just caught my eye that lithium reacts with nitrogen in a
different way than most metals. Seeing that it can produce both ammonia and nitrogen dioxide just with air and water. So I was seeing if anyone
heard of it. I was somewhat expecting a reply like "oh ya, it was used industrially as a catalyst in suchandsuch process, but it was too unreliable
or cost prohibitive to be economically viable."
To close the lithium loop I was thinking the LiOH would be electrolytically amalgamated to mercury or something.
Another thing I read was that a mercury cathode will strip the ammonium ion off a salt and (amalgamate it?) and release the NH3 and H2 from the
mercury once the current is off.
So, I also thought that maybe the mercury could hold the ammonia and the lithium as an amalgam, and they would react to make the amine. (or at least
hold them long enough to allow them to react somewhere else)
I am still curious about this though.
The reaction that makes the NO2 is stated as 'burning it' (with oxygen present). What is going on here? Is 'burning' required, or can it just be
oxidized to get it to release NO2? What chemistry goes on during 'burning'?
chemoleo, Li2O supposedly reacts with water to make LiOH, which I'm assuming the Li can me electrolytically amalgamated back into the mercury. And Li
left in dry air reacts with the nitrogen to make Li3N which then reacts with water to produce ammonia - so the ammonia is just made with air and
water. I'm not sure where the chlorine comes into play. BTW, what is a easy way to get NO2 gas (without using nitrates)
I am going to be building a NO2 generator by using corona discharge sometime in the near future. I thought maybe there was a more efficient way to
produce significant amounts of NO2/nitric from just air.
I don't take offense to your attitude (you may just be having a bad day), but you must notice my inexperience by the number of posts I have and my
questions. I am just looking for help in gaining more understanding about chemistry.kclo4 - 21-10-2008 at 16:45
The Easiest way I think would be either by corona discharge, or by oxidation of Ammonia. Ammonia is normally easy to get or make.
But to me it sounds like you are NO3- Hungry. As most begginers are; at least I was.
And I highly doubt the oxidation of Ammonia, or the Oxidation of Nitrogen is a very effective way of producing anything on the lab scale.
Other then buying nitrates as stump remover, fertilizer, Online, or at the pharmacy for something I believe you can extract it from organic material
such as manure. As I have recently stated in another forum post, I can't find the reference for this, and certainly wouldn't mind seeing it again. I
think this might be a way you can get nitrates - you'd have to look into it however.
As for producing NO2 with out using nitrates, I think you are stuck using air, or NH3 + Air.
From what I understand NO2 isn't all that useful other then for producing HNO3 on an industrial scale.kclo4 - 21-10-2008 at 16:54
If you used a source like Urea + Manure, it would probably help but honestly I haven't read much of it. I probably will, but I really just looked it
up for you.
This is something I've been wanting to try lately.
Edit: Magnesium can also burn in a Nitrogen Atmosphere to produce Magnesium Nitride, which reacts with water to produce Ammonia. - perhaps that also
interests you?
[Edited on 21-10-2008 by kclo4]OMG - 21-10-2008 at 17:03
Thanks for the links kclo4, I like the idea of getting the nitrates from the ground and letting the bacteria do the nitrifying for you. I might try
that out just for kicks.kclo4 - 21-10-2008 at 17:15
If you do, let me know how it goes. I'm interested in this and I haven't seen anyone have success or even make significant attempts at it.
Nitrate is getting harder and harder to get OTC, and that's how I like to buy it. Be good to have this as a method.
