# Tendency of formation of complexes in Group-1 elements

My book states the following:

Down the group(Group-1), tendency to form complexes decreases due to decrease in charge density$$\bigg(\frac{\text{charge}}{\text{radius}}\bigg).$$ This results in decrease in ion-dipole interaction between metal atom and the approaching ligand.

My logic-

There is a sudden availability of inner d-orbital while moving from 3rd period to 4th period. This can help accommodate the incoming lone pairs. So the tendency to form complexes must increase.

So combining the two facts, the trend should be irregular- decrease down till sodium, suddenly increase in potassium, and then slowly decrease till Caesium.

Can you provide any insight into the theory I have given above?

The fallacy is the assumption that "inner $$d$$ orbitals" become "available". Generally they are not, in the Group 1 and Group 2 metals. There are rare occasions where some evidence of $$d$$-orbital bonding is found for heavier G2 elements, but the impact is small and not widespread.
In Groups 1 and 2, where there are inner $$d$$ orbitals they are so well shielded that an electron placed there is not tightly bound and thus in general tends not to stay put. Somewhat ironically, $$d$$ orbitals become available to accept electrons from ligands just at the point where in neutral atoms they are being filled (i.e., the transition elements). This is because we add protons to the nucleus as we add electrons to the subshell, and the electrons we are adding to the $$d$$ orbitals do not fully shield the remaining $$d$$-orbital states from the increasing protic charge. It's as if the ligand electrons need to see their doppelgangers in the swimming pool before they figure out the water must be fine.
In truth, without readily available $$d$$ orbitals and with low charge density (except maybe lithium), the alkali metals are poor at forming complex ions. That includes complexing with hydroxide ions, thus explaining the unique basic strength of the hydroxides of these metals.