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Why are some anhydrous transition metal salts colored?

Here are some examples I found.

  • Anhydrous cobalt(II) chloride ($\ce{CoCl2}$), with sky-blue color.
  • Anhydrous chromium(III) chloride ($\ce{CrCl3}$), with purple color.
  • Anhydrous copper(II) chloride ($\ce{CuCl2}$), with yellow-brown color.

I already know that complex ion could undergo d-d transition from its unfilled d orbitals, and that's what makes them pretty-colored. (The wavelength of absorbed light corresponds to the gap in energies between d orbitals.)

But in anhydrous salt, which I think isn't a coordination compound, the d orbitals are degenerate. Thus, no energy gap and excitation could occur, no light is absorbed, and the compound would be colorless / white-colored as it reflects all wavelength.

I summarize my understanding in this problem as below.

  • Solution of complex ion could give color (with exception of $\ce{[Zn(H2O)6]^2+}$ and some other complex ions). e.g. $\ce{CuSO4 (aq)}$, which form $\ce{[Cu(H2O)6]^2+}$ complex ion with blue color.
  • Hydrated salt is indeed a coordination compound (found related question here), so it could give color. e.g. $\ce{CuSO4.5H2O (s)}$ with blue color.
  • Anhydrous salt, because it doesn't have any ligands coordinating with the metal ion, it's not a coordination compound. It's a mere ionic salt instead. Shouldn't it be colorless?

Is there any explanation to this case?

Note that in all compounds mentioned above, the metal ions are transition metal. We don't consider s-block metal ions such as $\ce{NaCl}$ and $\ce{MgSO4}$. (Additional question: They should be colorless, am I right?)

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    $\begingroup$ They are coordination compounds. $\endgroup$
    – Mithoron
    Commented Jan 24, 2016 at 17:39
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    $\begingroup$ The real question is why copper(II) sulphate is colourless in its water-free form. $\endgroup$
    – Jan
    Commented Jan 25, 2016 at 0:50
  • $\begingroup$ @Mithoron I thought ionic salt isn't called coordination compound by definition? $\endgroup$
    – Dean
    Commented Jan 25, 2016 at 17:42
  • $\begingroup$ Most so called salts are more covalent then ionic $\endgroup$
    – Mithoron
    Commented Jan 26, 2016 at 18:04

1 Answer 1

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A ‘mere ionic salt’, when inspected properly, turns out to be nothing different from a coordination compound. One side (the cation) is electron deficient while the other side (the anion) has spare electrons, so there will be some kind of coordination happening from one to the other. Indeed, all salts adopt crystal structures so that the favourable coordination interactions are maximised.

For anhydrous cobalt(II) chloride, I would immediately expect a closest packing with cobalt ions sitting in tetrahedral voids, occuping ¼ of them to give $\ce{[CoCl4]}$ tetrahedrons where each chloride coordinates two neighbouring cobalt ions. Analogously for chromium(III) chloride, I would expect chromium to occupy one-third of the anion structure’s octahedral voids, giving rise to $\ce{[CrCl6]}$ octahedrons where again each chloride coordinates two chromiums. Both are thus clearly coordination compounds. And going by what I know about $\ce{Cu^{II}}$, that salt will probably be a set of square-planar $\ce{[CuCl4]}$ fragments, again where each chloride coordinates two coppers.

Note that all of these are rough guesses (I wouldn’t even call them educated) since you can only predict solid-state (ionic) structures so far a priori. Oftentimes you can arrive at multiple possibilities and even more times you are surprised at what is actually adopted but then retrospectively you can usually explain the structure somehow with coordination.

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    $\begingroup$ Well, thanks for answering! I quite grasp what you're telling. Since they're kind of 'coordination' compounds, I'd suspect them to have some kind of excitation involving electrons from either the 'ligand' (anion) or the metal ion, and thus, their colors. $\endgroup$
    – Dean
    Commented Jan 25, 2016 at 17:40
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    $\begingroup$ Crystallography indicates that anhydrous cobalt chloride has cobalt in an octahedral environment with chloride anions as ligands. As the chloride is a very weak liagnd field ligand this accounts for the fact that cobalt chloride is blue and not red like hydrated cobalt chloride. $\endgroup$
    – Nuclear Chemist
    Commented May 1, 2018 at 6:07
  • $\begingroup$ @NuclearChemist But $[CoCl_6]^{4-}$ is faintly pink.? $\endgroup$
    – Rishi
    Commented Apr 28, 2021 at 13:42

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