In most simple compounds electrons occur paired in bonds, lone pairs, and core orbitals. However, the degree to which forming bonds lowers the molecules energy (in comparison to every atom just keeping its own electrons in a high-spin configuration and not forming bonds) is related to the degree to which the atomic valence orbitals overlap, from which the bonds would be formed.
The main reason why many transition metal compounds can keep unpaired electrons is that overlap of their d electrons to their ligands' s- and p-orbitals is not large enough to make forming actual covalent bonds energetically favorable over both the ligands and the metal just keeping their own electrons, rather than sharing them.