Question

I recently saw a lecture on hybridization and found two explainations for the different bond lengths in PCl5. One stated that equitorial bonds face less repulsion due to 120deg bonds but axial bonds face more repulsion due to its 90deg angle with the molecular plane.

The second one stated about the three equitorial hybrid orbitals being formed out of s, p_x and p_y orbitals and the axial bonds to be formed of the p_z and d_z^2 orbitals. Because the latter two are longer, the resulting hybrid orbital is also longer.

I have questions on the second explaination. How does this even work and when is this logic applicable? Is this actually a better model than the first explaination?

Answer 1

In the case of PCl5, the second explanation about hybrid orbitals is applicable. Phosphorus pentachloride (PCl5) forms five sp³d hybrid orbitals to account for the different bond lengths. This model of hybridization is a better fit for PCl5.

Step-by-step explanation:

In the case of PCl5, the second explanation about hybrid orbitals is applicable.

Phosphorus pentachloride (PCl5) has five P-Cl bonds directed towards the corners of a trigonal bipyramid. To form these bonds, the phosphorus atom hybridizes its 3s orbital, the three 3p orbitals, and one of the 3d orbitals to form five sp³d hybrid orbitals.

The equatorial bonds in PCl5 are formed by the overlapping of the three hybrid orbitals with the 3p orbitals and the s orbital. On the other hand, the axial bonds are formed by the overlapping of the 3d and 3p orbitals, which are longer and lead to longer bond lengths.

This model of hybridization explains the difference in bond lengths observed in PCl5 and is a better fit for this molecule.