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I've read in many books that paramagnetic substances are coloured while diamagnetic substances are colourless (yes there are some exceptions too). But nowhere i could find a reason for this phenomenon. I have an explanation so please tell me whether it is valid.

Paramagnetic substances have an unpaired electron. So the energy required to excite an unpaired electron is relatively less. Due to this the energy of visible light is sufficient for excitation and hence we are able to see the color when electron de-excites.

Please let me know if this is correct. Thank you.

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    $\begingroup$ The generalisation that paramagnetic compounds are coloured compared to diamagnetic ones is not a good generalisation. Consider, for example, that manganese sulphate (paramagnetic) is white to pale pink but potassium permanganate (diamagnetic) is an intense purple colour. $\endgroup$
    – matt_black
    Sep 21 at 11:23
  • $\begingroup$ True but in that case the colourisation is due for the color of respective anions, correct? In a situation where there is a colourless anion, is my explanation valid? $\endgroup$
    – Gnanadeep Sai
    Sep 21 at 12:00
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    $\begingroup$ There is IMHO rather very weak correlation than causation. Paramagnetic molecular entities have unpaired electrons, that MAY BE more inclined to absorb visible light due smaller difference of excited energy levels. It is far from being a rule. $\endgroup$
    – Poutnik
    Sep 21 at 12:16
  • $\begingroup$ @matt_black Additional example: the nice blue crystals of copper sulfate pentahydrate, $\ce{CuSO4 * 5 H2O}$ vs the white/gray powder of copper sulfate without any crystal water, $\ce{CuSO4}$. $\endgroup$
    – Buttonwood
    Sep 21 at 12:36
  • $\begingroup$ @GnanadeepSai I chose my example to illustrate that the colour of the ion containing the metal doesn't depend strongly on the magnetic state. The specific colour of a metal cation or anion varies a lot and doesn't correlate much with unpaired electrons. Also true for many non-metal compounds. Nitric oxide (NO) is not coloured despite being paramagnetic, as is oxygen (though very faintly coloured as a liquid). $\endgroup$
    – matt_black
    Sep 21 at 13:49

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As pointed out in the comments, this is not a good generalization.

You correctly stated that paramagnetic complexes have (at least) one unpaired electron, while diamagnetic ones have a closed shell. Your idea of colour is incorrect, what you describe is fluorescence. Generally speaking, we do not observe color because a compound emits light of a certain energy. We rather observe what is left from the light that shines on a compound. The absorbed part is the part we cannot see anymore. Whether a complex is coloured or not depends on the energy difference of the ground state and possible excited states, and whether a transition between these is spin-allowed or not. If there is a spin-allowed transition that has an energy in the range of visible light, we observe a colour. Consequently, there is no inherent connection between colour and magnetism.

I can, however, imagine how this coincidental correlation might arise: In complexes with an even number of d electrons, there is always the possibility to have a paramagnetic or a diamagnetic ground state configuration, e.g. the $d^8$ example in the diagram. When the orbital energy difference $\Delta\varepsilon$ is sufficiently small, the paramagnetic configuration will be the ground state. If it is large, the diamagnetic one is the ground state. The energy difference $\Delta\varepsilon$ also correlates with the energy of an electronic transition: the larger the difference, the shorter the wavelength we observe in a spectrum.

I assume that the energy difference at which the diamagnetic spin state becomes the ground state lies in the region where electronic transition have such high energies that their absorption does not happen in the visible region anymore and thus the complex appears colourless. This would justify what you found by connecting the electronic transition energy with the magnetism via an orbital energy difference.

I want to emphasize that an orbital energy difference is not the same as a state energy difference and serves purely as conceptual aid. Neither orbitals nor their energy differences are observables. d8 high spin and low spin configurations

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  • $\begingroup$ Okay so there is no relation with magnetism but we can conclude from the energy difference between orbitals? $\endgroup$
    – Gnanadeep Sai
    Sep 21 at 14:38
  • $\begingroup$ Yes, you can get an estimate of the colour of a complex from the orbital energy difference, if I correctly understood your comment. $\endgroup$
    – user137746
    Sep 21 at 14:53
  • $\begingroup$ I can't believe I'm the only other person to give your post a +1. In any event welcome to the community, you seem good at this, can you possibly attempt my question on colors? chemistry.stackexchange.com/questions/172258/… $\endgroup$
    – Neal Conroy
    Sep 24 at 0:05

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