I understand that the compounds of transition metal are colourful because of the appropriate energy gap between different d orbitals. But why aren't they also colourful in their solid pure state? For example, why isn't Fe colourful whereas, its compounds are?

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    $\begingroup$ The colour of TM compounds is typically because of electronic transitions. The electronic structure of bulk metal is not the same as the electronic structure of an isolated metal ion surrounded by ligands. Look up "band theory" for example. $\endgroup$ – orthocresol Aug 23 '20 at 17:54
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    $\begingroup$ @ Mathew. That is a good question. Do you know why copper is red ? And why Copper is an exception (apart from gold, which is due to relativist contraction) ? $\endgroup$ – Maurice Aug 23 '20 at 20:32
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    $\begingroup$ You said it already: it’s the metallic bonds that make things different. Do you expect a nitrogen atom and N2 to absorb light of the same wavelength? No, because the covalent bonds lead to a different set of orbitals with a different set of energies. Same with the metallic bonds. Note also that in a metal there are lots of metallic bonds, and each neighbour is further bonded to lots more atoms. In Cu(I) compounds there’s only one Cu-Cu bond and you can treat it as an isolated Cu2(2+) unit. $\endgroup$ – orthocresol Aug 24 '20 at 1:54
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    $\begingroup$ @orthocresol how is copper(I) dimeric when each ion has a closed subshell configuration? Zinc(I) is dimeric, but the ions bond to each other with electrons whose counterparts do not exist in copper/Group 11. $\endgroup$ – Oscar Lanzi Aug 24 '20 at 12:44
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    $\begingroup$ @OscarLanzi I am mortified. Was obviously too sleepy, probably thinking of Hg(I)....... Thanks for pointing out... $\endgroup$ – orthocresol Aug 24 '20 at 12:55

Colour is a property of electronic transitions: the transitions in metals are different to those in their compounds

In many discrete compounds of transition metals the colour arises because there are accessible electronic transitions between molecular orbitals in the molecules (usually involving metal d-orbitals) matching the energy of wavelengths of visible light. These transitions are often fairly narrow so giving a wide range of possible colours.

Pure metals are not like that. To simplify a lot, the definition of a metal is having a continuous band of electronic states (not a simple molecular orbital) for the electrons involved in conduction of electricity (the conduction band). This band involves the whole bulk substance, not just isolated atoms of the metal. So the transitions that might involve colour (d-d transitions in a molecule) are not present in the same way: instead transitions occur inside the conduction band or between orbitals and the conduction band. Either way, the narrow transitions giving colour in compounds are now much broader because a wide range of continuous electron energies exist in the conduction band.

This is a gross simplification, but it is good enough to explain the key observation if you don't want to get into the theory behind the electronic structure of metals.


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