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These two are simple ionic compounds , not complexes [if I do not consider water of hydration] then how can I rationalise the difference in colour?


marked as duplicate by Klaus-Dieter Warzecha, Community Mar 13 '15 at 6:51

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  • $\begingroup$ Cu3+ is still a different color than Cu2+, isn't it? $\endgroup$ – Lighthart Mar 13 '15 at 6:05
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    $\begingroup$ @Lighthart Honestly, I've never seen copper(III) in solution. Have you? It's not that copper(III) wouldn't exist, such as in ores, but as a naked, solvated cation? $\endgroup$ – Klaus-Dieter Warzecha Mar 13 '15 at 6:19
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    $\begingroup$ Do you really mean $\ce{CuCl3}$ or is it just a typo and you thought in $\ce{CuCl2}$? On a second note, copper(II)sulfate itself undergoes a colour change when converting the pentahydrate to the anhydrous form. $\endgroup$ – Klaus-Dieter Warzecha Mar 13 '15 at 6:37
  • $\begingroup$ LIGAND FIELD THEORY. Colors in these ions are based on electron transitions between non-degenerate d-orbitals. The energy of these transitions (thus the color) is a function of the metal oxidation state, the type of ligand/counterion, and the molecular geometry. Hydration or not d-orbitals are still going to distort based on those factors. $\endgroup$ – StevieD Mar 25 '16 at 6:32