In $PCl_5$, the phosphorus atom undergoes $sp^3d$ hybridisation. One $s$, three $p$, and one $d$ orbital mix to form five $sp^3d$ hybrid orbitals directed towards the five corners of a trigonal bipyramid. These orbitals overlap with the singly occupied $p$ orbitals of chlorine atoms to form five P–Cl sigma bonds.

Three P–Cl bonds lie in one plane, making an angle of $120^\circ$ with each other; these are termed equatorial bonds. The remaining two P–Cl bonds lie above and below the equatorial plane, making an angle of $90^\circ$ with the plane, and are called axial bonds.

The axial bonds are longer because the axial bond pairs suffer more repulsive interaction from the equatorial bond pairs. This makes them slightly longer and weaker than the equatorial bonds.