Sciencemadness Discussion Board - Reduction of NO3 salts

Reduction of NO3 salts

tom haggen - 21-4-2005 at 13:31

I really don't like the Idea of performing reactions at high temperatures due to the fact that I don't own an oxyacetlyene toarch, and a proper crucible. So here’s the question. With lithium being so readily available would it be possible to reduce NO3 salts with this akali metal?

(2Li+) + KNO3 ----> Li2O + KNO2

If this is possible what type of medium would you need for the reaction to take place in?

budullewraagh - 22-4-2005 at 06:37

if the lithium were as a cation, that wouldnt work. if the lithium were as the solid metal and you were to heat your reactants in an inert atmosphere, that would work. if this were to occur in anhydrous liquid ammonia, i think it would work as well.

tom haggen - 22-4-2005 at 11:47

Do you think that the NH3 would react with the NO3 salt to form NH4NO3?

CherrieBaby - 22-4-2005 at 13:20

The reduction of nitrates can be carried out by heat, light and ionizing radiation.[1]; addition of lead to fused NaNO3 at 400-450C [1]; reaction of nitrate in the presence of sodium ferrate and nitric oxide at ~ 400C [1]; contacting molten sodium nitrate with hydrogen [2]; and electrolytic reduction of sodium nitrate in a cell having a cation-exchange membrane, rhodium-plated titantium anode, and lead cathode [3]

Industrially sodium nitrite is prepared by contacting oxides of nitrogen with aqueous alkaline.[4]

[1] Mellor, Supplement to Mellor's Trestise on Inorganic and Theoretical Chemistry, vol VIII, suppl. II, part II, Wiley 1967
[2] USP 2294374
[3] Ger. Offen. 2940186
[4] USP 4009246

I'm told that you can use Iron instead of lead when heating fused nitrate salt [word of mouth only].

budullewraagh - 22-4-2005 at 14:38

"Do you think that the NH3 would react with the NO3 salt to form NH4NO3? "
no; NH4+ is significantly less active than any alkali metal

tom haggen - 22-4-2005 at 19:31

Well I might have to give this NH3 method a try. If anyone else has any opinions let me know.

unionised - 23-4-2005 at 05:01

Do you happen to have anhydrous liquid ammonia lying around somewhere waiting for you to think of a use for it?
I would be rather careful about mixing Li metal with any oxidant as powerful as KNO3, the mixture would be a reasonable explosive.

Mumbles - 23-4-2005 at 17:15

I think this is a fairly bad idea unless you have it under and inert atmosphere and keep it that way. I have a feeling that the reaction will not proceed as you think. If you look up the activity series, Lithium is higher than potassium. This means you will get a fun alkali metal thermite with free oxygen given off. Exciting, I know.

Use of just heat or lead would be best in my opinion. The reaction you have outlined will probaby just end explosion.

tom haggen - 23-4-2005 at 19:38

That’s good to know. I figured there had to be some kind of catch. Anyway, I've read a lot of the threads on reducing NO3 salts and IMO the method using LbO seems rather unreliable. If I had to go with any method I would probably choose the one that uses Aluminum as a reducing agent. But I don't currently have any means of heating Aluminum to its melting point. Anhydrous ammonia could be made by condensing ammonia gas in anhydrous conditions with dry ice and a proper solvent capable of reaching very low temperatures. However, I don't think this would be any easier than melting down some aluminum. Furthermore, I'm not interested in blowing myself up. I'm convinced that there has to be some better reducing agent out there than molten aluminum, so I will keep searching.

Marvin - 23-4-2005 at 21:17

I'm with Mumbles on this one, use lead (Pb) or carbon if you dont mind lower yeilds. Lead is only unreliable if the chemist is unreliable

Inconvenient, certainly.

Lithium might be a tad overkill as a reducing agent mainly due to the price, but it wont be quite as spectacular as described. You have to remeber the activity table is a lie, convenient to teach children but flawwed beyond the use of a general chemist.

tom haggen - 23-4-2005 at 22:00

I really don't have any experience with any of these reactions, but from what I have read the method using aluminum is far superior to the method using lead. Also, it is true that lithium is a little expensive but if you think about it 2 mols of lithium weighs what about 14 grams, and your one mol KNO2 you obtain is 1 mol that weighs about 102 grams. Sure lithium is expensive, but if you calculate a theoretical 100% yield, you’re getting about 102 grams products from 14 grams of reactants, that seems pretty darn effective.

Also I have been taking gen chem lately and haven't ever of the activity table. It sounds evil.

[Edited on 24-4-2005 by tom haggen]

unionised - 24-4-2005 at 01:13

C + 2KNO3 --> CO2 + 2KNO2
Less mass of carbon, compared to lithium(12 g per mole vs 7 but only 1 mole vs 2).
More product (2 moles not one)
and a reductant so cheap it's almost free.

Why piss about with neurotoxic, expensive lithium in impractical solvents for no proven yield when BBQ charcoal will do the job?

[Edited on 24-4-2005 by unionised]

frogfot - 24-4-2005 at 01:38

Does reduction with carbon really work? I tested it once with no results though I probably used wrong amounts.. Any details on the procedure?

Reduction with Al sounds like heating flash powder... not that fun..

I found some interesting procedures in a book. The most fun one involved heating aqueous solution of KNO3 and adding porous lead. The lead was prepared by precipitation from solution of lead (II) acetate by metallic zinc.
They're taking 2 parts lead to 1 part KNO3.
In the end of reaction Pb2+ is precipitated by bubbling CO2. This is followed by filtering, boiling down to consistency of butter, separation of precipitated KNO3 and addition of 90% ethanol untill KNO3 doesn't precipitate anymore. After filtering, soln is evaporated on water bath and dried in exicator.
This is said to give 60% yield of 80% KNO2 (the rest is KNO3)

Another procedure involved reaction between nitro ethyl and soln of NaOH in ethanol:

NaOH + C2H5ONO --> NaNO2 + C2H5OH

Yield is 91-96%. After recrystallisation from water product of purity 95,5-99,8% is obtained.

There's also some purification procedures..

Reference:
J.V.Karjakin, I.I.Angelov. Pure chemicals, 1974

EDIT:
Same book describes NaNO3/Pb method. It gives 75-85% yield. Seems like they're using pure lead, though they didn't mention any possibility of a runaway (the one I've mentioned in other nitrite thread).

[Edited on 24-4-2005 by frogfot]

Carbon is still the king.

Ramiel - 25-4-2005 at 00:05

I use the carbon reduction of potassium nitrate to get all of my needed nitrite. I've found that an excess of nitrate is needed to get a slurry that will maintain the reduction. When I separate the products through very careful re-crystallization, you get about 60% the expected yield of beautiful small square crystals first, then the familiar large needle-shaped nitrate crystals - both come of at quite distinct stages if one is careful.
- D

tom haggen - 25-4-2005 at 11:42

Just out of curiosity what kind of temperatures to you need in order for the reaction to proceed when using carbon??

Polverone - 25-4-2005 at 12:14

Ramiel, could you explain in a little more detail how you do your carbon reduction of KNO3? I've only tried this reaction a couple of times, and the resulting product produced copious red-orange fumes on addition of acid, but I did not attempt purification. I was under the impression that it would be difficult due to the very high aqueous solubility of KNO2. I am quite surprised that you speak of nitrite forming crystals before the nitrate. Do you use something other than water? As detailed in one of my old posts, it is also possible to produce KCN by the reaction of appropriate quantities of charcoal and KNO3 (yield undetermined). Who knew that humble potassium nitrate was so versatile?

N2o

ballzofsteel - 27-4-2005 at 04:25

Is there any chance of producing sodium nitrite by venting nitrous oxide(N2O) into a concentrated sodium carbonate sollution?
I know it works with NO NO2 etc but never a mention of N2O.

Just curious,thanks.

Esplosivo - 27-4-2005 at 05:21

N2O is a neutral oxide, and cannot therefore form salts with bases.

Ramiel - 27-4-2005 at 05:28

Wow, what a bother to get pure nitrite.. NO2? Though, if you wanted to be sure that your product was free of contaminants, even the Pb method wouldn't work I suspect.

Polverone: first with the purification, I definitely left out some interesting tidbits in the method, though I don’t know how much they affect results!
A number of acids were added(Inorg. and Organic), and some marginal results were recorded by re-crystallization with trace amounts of HCl in the solution. It was noted that the crystals formed were almost all perfectly cubic and regularly sized. It was probably just an anomalous result; however I was prepared to move on a hunch. Some more of the same solid was solved in distilled-deionised water this time, (having been properly filtered et cetera

. I added some KCl in trace amounts (another hunch) and the results were, well, useful! As stated before, a distinct level of re-crystallization was observed.
I hope this helps some.

As for the method? I use a coffee can on top of a steel plate, heated by a little fire out the back, lo-tech, and the hot spots help

(joke! LOLOLOL!!!1)

Given:
2.KNO3 (L) + C(S) ==> KNO2 (L).(S) + CO2 (G)
and:
KNO2 (L) + C(S) ==> useless crap clogging up the reaction(S)
(Can you tell I love the function?

)

You need an excess of KNO3 for two reasons - so you don't get further reduction of the pot.salt to useless crap, and also so the excess pot nitrate makes a slurry that is conducive to reduction - in a stoichiometric amount, the slurry is too thick, it doesn’t melt thoroughly, and the reaction stalls. If one has fair temperature control then;
First heat the whole thing to a liquid slurry, 3 moles pot.nitrate to 1 mole carbon,
Initiate it by dropping a teeny bit of lit sugar/pot.nitrate in.
The reaction goes ahead quite nicely if you do it properly. No stalling, no toxic fumes, just a bit of spluttering which can be solved with a loose cover.

The only problem with this is you will probably have to sacrifice the coffee can – the slurry adheres amazingly well to the metal as it cools!

I can see myself banning this thread, so heed this as a warning you Polverone and Ramiel persons! no more practical discussions of pyrotechney.

- D