The salient feature of the theory are summarised below.
(i) The central metal ion has a number of empty orbitals for accommodating electrons donated by the ligands. The number of empty orbitals is equal to the coordination number of the metal ion for the particular complex.
(ii) The atomic orbitals (s, p or d) of the metal ion hybridize to form hybrid orbitals with definite directional properties. These hybrid orbitals now overlap with the ligand orbitals to form strong chemical bonds.
(iii) The d-orbitals involved in the hybridization may be either inner (n - 1) d orbitals or outern d-orbitals. The complexes formed in these two ways are referred to low spin and high spin complexes, respectively.
(iv) Each ligand contains a lone pair of electrons.
(v) A covalent-bond is formed by the overlap of a vacant hybridized metal orbital and a filled orbital of the ligand. The bond is also sometimes called as a coordinate bond.
(vi) If the complex contains unpaired electrons, it is paramagnetic in nature, while if it does not contain unpaired electrons, it is diamagnetic in nature.
(vii) The number of unpaired electrons in the complex points out the geometry of the complex and vice-versa. In practice, the number of unpaired electrons in a complex is found from magnetic moment measurements as given in Table.
Relation between unpaired electron and magnetic moment
| Magnetic moment (Bohr magnetons) |
0 |
1.73 |
2.83 |
3.87 |
4.90 |
5.92 |
| Number of unpaired electrons |
0 |
1 |
2 |
3 |
4 |
5 |
Thus the knowledge of the magnetic moment can be of great help in ascertaining the type of complex.
(viii) Under the influence of a strong ligand, the electrons can be forced to pair up against the Hund's rule of maximum multiplicity.
