[bradvincent]

Take the equation for conversion of free nitrogen and oxygen to nitrate and vice versa:

2NO₃ N₂ + 3O₂

Does this converting of free nitrogen and oxygen to nitrate release or consume energy?  Does this converting nitrate to of free nitrogen and oxygen release or consume energy? 

I remember vaguely counting the types of bonds formed on each side to solve this, but am very rusty.

FYI - I am not building anything that explodes

[Corribus]

Use heats of formation to determine if the reaction is exothermic for a given temperature.  Since two of those are already standard reference compounds, you only have to look up one.

(By the way, "release or consume energy" is different from "release or consume heat".  Make sure you are after what you think you're after. From the context I'm assuming you want the latter, not the former as specified.)

[bradvincent]

I was unaware "release or consume energy" is different from "release or consume heat", at least in any reaction that does not release significant energy that is not heat (energy as light, sound, etc).

Chemistry.about.com gives me -206.6 kJ/mol heat of formation for nitrate.  I am guessing the other 2 are what you meant by "standard reference compounds", and that means they have a heat of formation of 0 (by definition) and the reaction release 206.6 kJ per mole of nitrate formed, or consumes 206.6 kJ going the other way.

I was unaware I provided a context.

[Corribus]

First, when we speak of energy changes for a reaction, this can mean many things but usually it refers to changes in (Gibbs) free energy, which involve both entropic and enthalpic differences between reactants and products. Heat consumption or release primarily refers to the enthalpic portion. The "context" I alluded to was your mention of using bond energies. This is also a viable way of estimating heat loss/gain by a reaction, but it is more cumbersome and less precise in this case because bond energies do not take into consideration of intermolecular interactions, which are incorporated in widely-available heat of formation data.

Second, the way you determined the enthalpy change is correct.  However do note that your equation is not balanced for charge and therefore your decomposition reaction is not realistic.  Nitrates are anions in aqueous solution and always have a counterion which needs to be incorporated into the equation in some way.

A more realistic equation might be something as follows:

4NO₃^-(aq) 2O^2-(aq) + 4NO₂(g) + O₂(g)

Or some such...
(Likely the superoxide will be bound to one of your metallic cations and precipitate as a solid; either that or it will disproportionate with water into hydroxide and oxygen gas, so even the way I have formulated it here is probably unrealistic, but I think you get the idea...)

[curiouscat]

First, when we speak of energy changes for a reaction, this can mean many things but usually it refers to changes in (Gibbs) free energy, which involve both entropic and enthalpic differences between reactants and products.

Probably, besides the point, but I'm curious whether this is how everyone interprets "energy changes"? Coz' I would by default assume an enthalpy change if nothing were mentioned. Worst case, in my DFT days, I'd have assumed an Internal Energy change.

But somehow I'd never think to assume a Gibbs free energy change.

Just curious what other people's defaults are.

[curiouscat]

(By the way, "release or consume energy" is different from "release or consume heat".  Make sure you are after what you think you're after. From the context I'm assuming you want the latter, not the former as specified.)

Another tangent, but this got me thinking: What is the maximum fraction of the enthalpy change of any one reaction that people know that gets liberated as non-heat modes?

Are there reactions which convert to light (or I don't know what, sound?) close to 100% of their enthalpy deficit?

Just curious.

[Corribus]

@curious

I would (and did, here) usually interpret it as enthalpy change as well, and maybe I'm too pedantic when I gripe about it, but I think it's important for people to recognize that energy does not always equal heat. Speaking of an energy change, in other words, can be ambiguous.  It's always better to state exactly what you mean to limit confusion.

There are of course reactions that generate a considerable amount of light. A reaction which converts all of its energy change into light would require a chromophore with exactly 100% fluorescence yield. There are some that come close, but none to my knowledge whose excited states are the products of a chemical reaction. I'm not sure what fraction of, for example, luminol's enthalpy change is manifested as light (vs. heat) when it is converted to its downstream product by peroxide. I'm sure this information is available somewhere.

[curiouscat]

  It's always better to state exactly what you mean to limit confusion.

Yes, agreed!

[Corribus]

(Likely the superoxide will be bound to one of your metallic cations and precipitate as a solid; either that or it will disproportionate with water into hydroxide and oxygen gas, so even the way I have formulated it here is probably unrealistic, but I think you get the idea...)

Just wanted to correct this. No superoxide here, just oxide ion, which is similarly unstable in water and will form hydroxide ions... but no oxygen gas.