Why are "s" + "p" overlaps directional bonds?

Because by definition, covalent bonds are directional in nature.

Covalent bonds require orbital overlap, that is, the sharing of valence electrons to some extent. And orbitals can only overlap in particular orientations (while still being orthogonal to each other) to form these bonds. So the molecules that form have fairly definite shapes. This is also said pretty well here.

The ultimate goal is stability at the bottom of the potential well (at #r = r_(eq)#), and that is accomplished by balancing:

• the repulsive force between the two nuclei (#A//B#)
• the attractive force between the nucleus of atom #A# with electron density from atom #B# and between the nucleus of atom #B# with the electron density from atom #A#.

The former tends to infinity as the internuclear distance #r# approaches #0#, where #0 < r < r_(eq)#.

The latter approaches energy zero from below and from the left at #r_(eq) < r < oo#:

Two atoms in space can only make a proper covalent bond by approaching each other in a straight line, and thus, the covalent bond-forming process is directional.

Are #s# and #d# overlaps directional?