Nitrogen dioxide (NO2) is a reddish brown molecule, while its dimer nitrogen tetroxide (N2O4) is colorless. Can you give a reasonably detailed explanation for this from electronic structure theory?

Previous queries were not answered because the question posed was deemed to be too broad.

  • $\begingroup$ I'm afraid that like the previous question, this will also be deemed as too broad very soon. EST is a very broad and complicated theory and may be hard to explain that in a single answer. $\endgroup$ Nov 20, 2017 at 4:27
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    $\begingroup$ No, I think what made the other question too broad is wanting to include the other coloured gases, which would require an explanation for every single gas. $\endgroup$
    – Jan
    Nov 20, 2017 at 11:24
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    $\begingroup$ You are asking why NO2 has low lying energy levels compared to its dimer. (It just happens that these low levels lie in the visible part of the spectrum). To put the question the other way round, why would you expect the energy levels to be similar when the molecules are different? Perhaps more interesting is to ask why N2O4 is planar yet has a very long 0.175 nm N-N bond whereas the NO bond lengths and ONO angles are similar to NO2. $\endgroup$
    – porphyrin
    Nov 21, 2017 at 11:07
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    $\begingroup$ The color of NO2 is due to presence of single unpaired electeon, while in its dimer N2O4 that lone electron got paired which cause loss of its color. $\endgroup$
    – Deadwing
    Mar 9, 2021 at 19:18
  • $\begingroup$ Unanswered duplicate: 1 $\endgroup$ Mar 20 at 13:48

1 Answer 1


To understand this, we must take a look at these molecules one by one.

  1. Nitrogen Dioxide[NO2]: enter image description here

    This here is its resonance structure/canonical structure.

NO2 is a molecule in which a Nitrogen is connected with two atoms of Oxygen and its shape is Bent. Total number of valence electrons is 17 with each Oxygen atoms contributing 6 and 5 electrons are from Nitrogen.

According to MOT (Molecular Orbital Theory), NO2 has an unpaired electron in π* (Antibonding) orbital due to which, orbitals of Nitrogen and Oxygen undergo 2p-2p overlap. Meanwhile this unpaired electrons makes our molecule a paramagnetic one.

This unpaired electron in π* (Antibonding) orbital can react with any external electromagnetic radiations and guess what, visible light is also an electromagnetic radiation. The reaction between this unpaired electron and visible light causes electron transition.

enter image description here

When electrons goes to higher energy orbital from lower energy orbitals some form of energy i.e LIGHT is absorbed, if the electron goes vice-versa then energy (LIGHT) is released.

The presence of unpaired electrons in paramagnetic specie creates molecular orbitals that are partially filled or has unpaired spins.these electrons occupy higher energy orbitals which are more prone to electronic transitions. So, paramagnetic specie are more viable to absorb light at lower frequency as they have longer wavelength. Hence, NO2 absorbs lower frequency light in the visible region of the spectrum, specifically in the blue-green range.

As a result, NO2 appears reddish-brown in colour.

2.Dinitrogen Tetroxide(N2O4): enter image description here

 This here is its resonance structure/canonical structure. 

THIS is a dimer of NO2 consisting of two nitrogens and four oxygens. In this molecule due to dimerisation, the unpaired electrons in the π*(antibonding) orbitals of each NO2 molecule pair up together to form a stable, diamagnetic molecule.

Since it has no unpaired electrons, No interaction with external electromagnetic waves and so, no absorption of light in the visible region. Hence, it appears colourless.

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    $\begingroup$ Close to a nice answer,but I think this needs a bit more - why does a molecule being paramagnetic make it able to absorb light at comparatively low frequency? If paramagnetism is the cause of colour why is NO colourless? $\endgroup$
    – Ian Bush
    Mar 20 at 9:50

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