The non-stoichiometric form of nickel (II) oxide and copper (II) oxide are black, with the same square-planar (at metal) solid-state structure, (with crystal defects), so the metal centres are bridged by oxides. The assignment of oxidation state is +II; therefore, could consider a fragment of the structure as: -O-Cu(II)-O-Cu(II)-; but, with redox-active metals, this also:
-O-Cu(III)-O-Cu(I)- thus. Electrons can drift through the structure giving rise to intervalence electronic transitions, with strong involvement of oxygen atoms. The influence of pie-donating oxygen prevents the formation on Cu(0), the native metal.
Intervalence transitions can give rise to intense colours (e.g. black); normally, forbidden to colours arising from d-d electron transitions. (This, also seen with Prussian Blue, where Fe(II) & Fe(III) are bridged by cyanides.) A discrete Cu(II) centre would have a weak, less-intense, colour, arising from said d-d transitions.
Stabilisation of highly-polarising Cu(III) & Ni(III) is achieved by the oxide, which is a very electronegative and pi-donating ion, stimulating low activation-energy electron-flow.
Transitions between p and d orbitals (as in permanganate: metal and oxide) give rise to very intense colours. These extended oxide lattices are also susceptible to being defective, so missing ions; or, ions in the wrong location in the lattice. This also gives rise to colour. An example of this is seen in zinc (II) oxide, when heated. In this case it is not mixed valence; but, loss of oxygen from the lattice, that gives rise to defects that allow lower energy (than in "complete" ZnO) electronic transitions to occur and absorption at the UV-end of the visible region. On cooling the lattice picks up oxygen, from the air, becoming complete again, reverting to its original white from the, high-temperature, yellow colour.