treasurefish - 16-8-2006 at 12:41
With regards to the “Cold-Pack” endothermic reaction:
NH4NO3 (s) -> NH4+(aq) + NO3-(aq)
H = +28.1 kJ/mol
Why doesn’t the reaction continue to the following step and instead produce tons of heat?
5NH4+ + 3NO3- -> 4N2 + 9H2O + 2H+
H = -1483 kJ/mol
I saw this reaction listed in the review notes for a chemistry course at: http://www.howard.edu/ceacs/Departments/Civil/cglass/Env.Mic...
I'm thinking that since water molecules are polar, there is a strong attraction of water molecules to both cations and anions through the
charge-dipole intermolecular interaction. Doing a review on the internet (keyword: anammox), it seems that sub-sea bacteria perform this
denitrification function, but is there any way to make this happen chemically?
If a catalyst or the right activation temperature was determined, then the reaction should continue without any further additional inputs due to the
exothermic potential. Right? Is there a catalyst that can get this started? What activation energy would be required?
Thank you.
12AX7 - 16-8-2006 at 14:06
For one, the reaction isn't stoichiometric with NH4+.NO3-, which is 1:1, not 5:3. Even so, the 3/5ths that would react (oh, and the two H+ left would
end up as HNO3, since 2NO3- wouldn't be used) as such would rather react as N2H4O3 > N2O + 2H2O, which occurs something above the melting point of
AN.
Biology can assemble and dissassemble N2, NH4+, NO2- and NO3-, but water alone can't do it. Bacteria use very specialized pathways to perform these
reactions (and thanks to evolution, they do it very efficiently), but in bulk, you aren't going to get either until you heat the heck out of the AN,
or detonate it (not easily done).
In short, it takes too much energy to break up the stable, symmetrical NH4+ and NO3- molecules so the O's can combine with the H's and the N's with
each other.
And for that matter, you ought to ask why AN doesn't spontaneously decompose on its own -- there's no reason water need be involved in your combustion
reaction; it's an ionic solid, and indeed the NH4+ and NO3- groups are closer together than they ever are in solution.
Tim
chemoleo - 16-8-2006 at 14:35
Well the reverse reaction takes place in the Haber Bosch process:
NH3 + H2O + O2 --> NO2, NO + NH3 + O2 + H2O --> NH4NO3 etc .
The catalyst is Pt (hot), and if you run the reaction with surplus NH3, you do indeed get ammonium nitrate. The reaction is even self-sustaining,
doesn't require outside heat.
The activation energy is what separates the reactants from the products. Pt just lowers it.
Anyway, your reaction is no different to asking, why doesn't sugar C6H12O6 instantaneously combust, in fact why not most organic oxyen-containing
compounds? They don't again due to the activation energy. At high temp, in a closed container, sugar will decompose, essentially burn with itself to
produce water, carbon, CO , and tar products.
treasurefish - 16-8-2006 at 18:31
12AX7 and Chemoleo - thank you for your time - I appreciate your patience with me.
I am eager to find a reaction that uses cheap reactants to generate lots of heat and N2, and maybe some H2O and H+. Basically hot denitrification. It
would be best if it did not involve metals and did not create any other "waste" products. It doesn't matter if it is "high energy," but having this
quality wouldn't hurt either.
My next idea is: NH4+ + NO2- → N2 + 2H2O
The reaction rate should be fast (2.7 x 10-4/s x [NH4+][NO2-]), exothermic (-107 kcal/mol), and spontaneous (-112 kcal/mol). Seems like the best of
worlds to me - how can I find out what the necessary conditions are to make it happen? Any experience out there?
Since NO2- is perishable in the form of nitrous acid, I've read that NaNO2 is a good vehicle. If so, is there a good companion for NH4 as well?
I'm sure these questions are very simple for a lot of folks out there, but this is truly the only forum where the simple-minded are rewarded with
genious responses. Sounds like a good ad - heh?
[Edited on 17-8-2006 by treasurefish]
[Edited on 17-8-2006 by treasurefish]
Nicodem - 16-8-2006 at 20:46
That reaction was used as a simple laboratory source for N2. If I remember corectly a solution of NaNO2 was slowly introduced with a addition funnel
to a solution of NH4Cl in the flask (at room T). The evolved N2 was then dried and used for whatever aplication.
You should be able to find details in Vogel's or some other books on practical chemistry. But I don't think the reaction mixture heats much if at all.
The enthalpy you provided is quite low and you should calculate the free energy (deltaG) to see if any heat is produced at all at the room
temperature.
not_important - 17-8-2006 at 02:02
For a more extreme case of activation energy getting in the way, consider this:
NH4(+) => {some step we're going to ignore} => He + N(+) + 26.7 MeV
This would give you hot nitrogen 8-)
Because of activation energy we have a solid planet instead of a cloud of hot, dirty helium. It also means that it really hard to threaten a city
with distruction by means of a bottle of mineral water.
More seriously, how hot is 'hot'? Not creating much of anything but N2 would be tricky, water is the obvious byproduct to be expected - how much of it
is 'OK'?
Here's a bit more detailed ammonium nitrite reaction
http://www.uq.edu.au/_School_Science_Lessons/UNChemLesson.ht...
http://www.owlnet.rice.edu/~chem122/Lab122/propgases/index.h...
Thank you Nicodem and Not_Important!
treasurefish - 18-8-2006 at 12:00
You guys set me off on the right direction!
Take care.