# Oxygen production from thermal decomposition of metal oxides

Does thermal decomposition of metal oxides always produce oxygen?

In addition to the examples given by Jannis Andreska, metal peroxides, such as barium peroxide and lithium peroxide also release oxygen upon heating:

$$\ce{2 BaO2 ->[\text{T = 700 °C}] 2 BaO + O2} \\ \ce{2 Li2O2 ->[\text{T = 195 °C}] 2 Li2O + O2}$$

This was once used to produce pure oxygen.

On the other hand, some, if not most, metal oxides are extremely heat resistant. For example, once sintered at temperatures between $1700$ and $\pu{2000°C}$, magnesium oxide $(\ce{MgO})$ can be heated up to its melting point ($\sim\pu{2800°C}$) without decomposition and can thus be used as a lining for heat sensors.

A lot of metal oxides decompose under liberation of $\ce{O2}$ when subjected to sufficiently high temperatures. For example, $\ce{HgO}$ decomposes at temperatures above $\pu{500^\circ C}$ into elemental mercury and oxygen:

$$\ce{2HgO ->~ 2Hg + O2}$$

However, disproportionation is another possible reaction pathway. An example is the thermal decomposition of iron(II) oxide to elemental iron and iron(II, III) oxide at temperatures below $\pu{575^\circ C}$ ($\ce{FeO}$ is more stable above that temperature; sources: 1, 2):

$$\ce{4FeO \rightleftharpoons Fe + Fe3O4}$$

Another example is the decompostion of $\ce{SnO}$ at temperatures between $700\mathrm -\pu{1000K}$, which proceeds via an intermediate $\ce{Sn3O4}$ mixed oxide. Tin disproportionates from the oxidation state +II in $\ce{SnO}$ to $0$ in $\ce{Sn}$ and +IV in $\ce{SnO2}$ (1, 3).

\begin{aligned} \ce{4SnO ->~& Sn3O4 + Sn} \\ \ce{Sn3O4 ->~& 2SnO2 + Sn} \end{aligned}