Why do transition metals form particularly stable complexes?
To illustrate my answer, I am going to use a modified version of the picture of a simplified molecular orbitals scheme Professor Klüfers (LMU Munich) uses in his internet scriptum to the coordination chemistry I course.



On the left you can see the metal’s orbitals, on the right the group orbitals of six $\unicode[Times]{x3C3}$ donors (assuming no $\unicode[Times]{x3c0}$ interactions of any kind). $\mathrm{e_{g}}$ and $\mathrm{t_{2g}}$ correspond to the metal’s d-orbitals, $\mathrm{a_{1g}}$ is the 4s orbital and $\mathrm{t_{1u}}$ are the 4p-orbitals.

You can clearly see that the 3d-orbitals are similar in energy to the ligand group orbitals, and that they have the correct symmetry to mix. The greatest stabilisation comes from the $\mathrm{e_g}$ orbitals.

Empty d-orbitals can thus be thought to help with coordination by stabilising the coordination compound. Completely filled d-orbitals (e.g. $\ce{Zn^2+}$) are contraproductive and therefore zinc’s coordination chemistry is much like sodium’s.