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This is basically a continuation of the question-"Why is potassium monoxide (K2O) coloured?"

I knew that color of alkali metal oxides deepens down the group:

  • Lithium oxide ($\ce{Li2O}$) is the lightest alkali metal oxide and a white solid. It melts at 1570 °C.
  • Sodium oxide ($\ce{Na2O}$) is a white solid that melts at 1132 °C and decomposes at 1950 °C.
  • Potassium oxide ($\ce{K2O}$) is a pale yellow solid that decomposes at 350°C.
  • Rubidium oxide ($\ce{Rb2O}$) is a yellow solid that melts at 500 °C.
  • Caesium oxide ($\ce{Cs2O}$) is a yellow-orange solid that melts at 490°C.

The reason is due to what @Oscar said in the answer of previous question:

If we go further down the alkali metal oxides the band gap gets smaller, and more visible light with longer blue/green wavelengths may be absorbed. Then the color becomes stronger and more reddish.

But, it is observed that not only color of normal oxides deepens down the group but the peroxides and superoxides also follow the trend:

  • Lithium peroxide ($\ce{Li2O2}$) is a white solid that melts at 195 °C.
  • Sodium peroxide ($\ce{Na2O2}$) is a pale yellow solid that melts at 460 °Cand boils at 657 °C.
  • Potassium peroxide ($\ce{K2O2}$) is a yellow solid that melts at 490 °C.

  • Sodium superoxide ($\ce{NaO2}$) is a yellow-orange solid that melts at 551.7°C.
  • Potassium superoxide ($\ce{KO2}$) is a yellow solid that decomposes at 560°C.

  • Lithium ozonide ($\ce{LiO3}$) is a red solid which is produced from caesium ozonide via an ion-exchange process.
  • Sodium ozonide ($\ce{NaO3}$) is a red solid which is produced from caesium ozonide via an ion-exchange process.
  • Potassium ozonide ($\ce{KO3}$) is a dark red solid which is produced when potassium is burned in ozone or exposed to air for years.
  • Rubidium ozonide ($\ce{RbO3}$) is a dark red solid which is produced when rubidium is burned in ozone.
  • Caesium ozonide ($\ce{CsO3}$) is a dark red solid which is produced when caesium is burned in ozone.

What is the reason behind the deepening of the color? Is it due to the same reason as @Oscar pointed out? Or is it any other reason? Although superoxides and ozonides are paramagnetic and are colored, the peroxides are diamagnetic but still they are colored. Why?

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  • $\begingroup$ Are you sure that ozonides are diamagnetic? $\endgroup$ – permeakra Apr 9 '17 at 13:29
  • $\begingroup$ @permeakra, yes, my bad, rectified. $\endgroup$ – Nilay Ghosh Apr 9 '17 at 13:39
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Most likely, we are seeing the same band structure effects.

There is an additional trend, whereby the compounds become darker and more reddish as one adds more oxygen. In oxygen-chain species such as $\ce {O2^{2-}} $, antibonding molecular orbital are occupied. So The valence band containg these electrons is raised. That lowers the band gap enabling absorption of more visible light.

Incidentally, the visual effects described here are limited to alkali metal (when oxygen is the electronegative element). When we move to the corresponding alkaline earth oxide the monoxides are white and the peroxides ($\ce {Mg}$ and heavier) are white too. Less strongly ionic bonding and a lower cation/anions ratio lead to less negative charge on the oxygen species, thus lower valence band energies and gaps remaining large enough to remain uv-absorptive only.

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