# Do all metal nitrates give nitrogen dioxide on decomposition?

Some metal nitrates decompose to give nitrites: $$\ce{2MNO3 -> 2MNO2 + O2}$$
whereas some metal nitrates decompose to give $$\ce{NO2}$$. A possible reaction is $$\ce{M(NO3)2 -> MO + NO2 + O2}$$

Based on my observation my question is that, is it true that all metals (including transition metal) nitrates decompose to $$\ce{NO2}$$ apart from $$\ce{[Na , K, Rb, Cs]}$$ (alkali metals which give nitrites instead) ?

• $\ce{NO2}$ is pretty much always there when a nitrate is decomposed (also for $\ce{MNO3}$). The question is, how much. And please try not to use chemical formulas in titles (especially those requiring MathJax), use chemical name instead. Apr 23 at 12:29
• Strongly related: chemistry.stackexchange.com/questions/51769/… Apr 23 at 12:40
• Apr 23 at 12:41
• Thanks @NilayGhosh I just read all of your suggestions. They did answer my second potential question of whether the metals would exits in free state or as oxides after decomposition. Although for this question I was seeking a possible generalization and affirmation on this observation and to know any exceptions if they might exist. Apr 23 at 12:54

TL;DR: Decomposition of nitrites always afford $$\ce{NO2}$$ if you apply the correct reaction condition to it.

To generalize the decomposition of metal nitrates (I concluded from the article above, and hereby using a generic monovalent metal M for the equations), we can say this:

• The first stage of decomposition is always the decomposition to nitrites: $$\ce{MNO3 -> MNO2 + 1/2 O2}$$

• Note that nitrites are almost always (if any) less thermally stable than the corresponding nitrates. Thus, the formed nitrites from the reaction above quickly decomposes. I think that (and from the article), nitrites decompose to form the oxide and $$\ce{N2O3}$$, the latter disproportionates into $$\ce{NO}$$ and $$\ce{NO2}$$: $$\ce{2 MNO2 -> M2O + NO + NO2}$$.

• Everything becomes interesting from here. It is due to the fact that the nitrite formed from the first step and not decomposed (kinetics matter) can be oxidized by either the $$\ce{NO2}$$ or $$\ce{O2}$$ back to the nitrate. This is especially significant when the nitrite is stable enough ($$\ce{KNO2, NaNO2}$$ etc). In fact, the decomposition of $$\ce{KNO3}$$ is reversible (equilibrium) at about 650 - 750 °C in the atmosphere. In the other hand, nitrites of transition metals (which are mostly covalent in character) are so unstable that they decompose rapidly and thus only appear as the so-called reactive intermediates.

• Finally, the oxide can decompose into the metal and $$\ce{O2}$$. This is a characteristic of metals that lie in the less-reactive end of the reactivity series. If the decomposition temperatures of the oxide and the nitrate are different enough, you can control the temperature to get the oxide or the metal ($$\ce{HgO}$$ from $$\ce{Hg(NO3)2}$$ is an example.)

The re-oxidation of stable nitrites can be avoided by two ways: either do the reaction under a flow of inert gas to displace the gaseous products away, or crank up the heat more to increase the reaction temperature, which means faster decomposition of nitrites. For example, decomposition of $$\ce{KNO3}$$ goes to completion at about 800 °C.

Reference

1. High Temperature Properties and Decomposition of Inorganic Salts Part 3, Nitrates and Nitrites, Kurt H. Stern, Journal of Physical and Chemical Reference Data 1, 747 (1972); DOI: 10.1063/1.3253104