Chemistry - Coordination Chemistry Question with Solution | TestHub

When a transition metal ion (usually) is involved in octahedral complex formation, the five degenerate d orbitals split into two sets of degenerate orbitals (3 + 2). These are three degenerate orbitals of lower energy (${\mathrm{d}}_{xy}$, ${\mathrm{d}}_{yz}$, ${\mathrm{d}}_{xz}$) and a set of degenerate orbitals of higher energy (${\mathrm{d}}_{x^2-y^2}$, ${\mathrm{d}}_{z^2}$). The orbitals with lower energy are called ${\mathrm{t}}_{2\mathrm{g}}$ orbitals, and those with higher energy are called ${\mathrm{e}}_{\mathrm{g}}$ orbitals.

In octahedral complexes, a positive metal ion may be considered to be present at the center and negative ligands at the corners of a regular octahedron. As the lobes of ${\mathrm{d}}_{x^2-y^2}$ and ${\mathrm{d}}_{z^2}$ lie along the axes (i.e., along the ligands), the repulsions are greater, resulting in higher energy. The lobes of the remaining three d-orbitals (${\mathrm{d}}_{xy}$, ${\mathrm{d}}_{yz}$, ${\mathrm{d}}_{xz}$) lie between the axes (i.e., between the ligands). The repulsion between them is less, resulting in lower energy. In octahedral complexes, if a metal ion has more than 3 electrons, then for pairing them, the options are:

(i) Pairing may start with the 4ᵗʰ electron in ${\mathrm{t}}_{2\mathrm{g}}$ orbitals.

(ii) Pairing may start normally with 6 electrons when ${\mathrm{t}}_{2\mathrm{g}}$ and ${\mathrm{e}}_{\mathrm{g}}$ orbitals are singly filled.