As the oxidation state of the metal increases, so also does the amount of splitting of the d orbitals.
Changes of oxidation state therefore change the colour of the light absorbed, and so the colour of the light you see.


Taking another example from chromium chemistry involving only a change of oxidation state (from +2 to +3):

The 2+ ion is almost the same colour as the hexaaquacopper(II) ion, and the 3+ ion is the hard-to-describe violet-blue-gey colour.

Splitting is greater if the ion is octahedral than if it is tetrahedral, and therefore the colour will change with a change of co-ordination. Unfortunately, I can't think of a single simple example to illustrate this with!

The problem is that an ion will normally only change co-ordination if you change the ligand - and changing the ligand will change the colour as well. You can't isolate out the effect of the co-ordination change.

For example, a commonly quoted case comes from cobalt(II) chemistry, with the ions [Co(H2O)6]2+ and [CoCl4]2-.

The difference in the colours is going to be a combination of the effect of the change of ligand, and the change of the number of ligands.