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Does thermal decomposition of metal oxides always produce oxygen?

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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.

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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}

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