Formation of Molecular Orbitals — Definition

Definition

Imagine atoms as having specific 'rooms' (atomic orbitals) where their electrons reside. When two atoms decide to form a molecule, their 'rooms' don't just stay separate; instead, they merge or combine to create new, larger 'rooms' that belong to the entire molecule. These new, molecule-wide rooms are called molecular orbitals (MOs). The process of forming these MOs is governed by a principle called the Linear Combination of Atomic Orbitals (LCAO).

Think of it like this: each atomic orbital is described by a mathematical wave function, which tells us about the probability of finding an electron in a particular region around the nucleus. When two atoms come close enough to bond, their atomic orbitals' wave functions can interact in two main ways, similar to how waves in water can interact.

1
Constructive Interference: — If the waves combine 'in phase' (like two crests meeting), they reinforce each other. This leads to an increased electron density between the nuclei, pulling them closer. This type of interaction forms a bonding molecular orbital (BMO). Electrons in BMOs stabilize the molecule and have lower energy than the original atomic orbitals.

1
Destructive Interference: — If the waves combine 'out of phase' (like a crest meeting a trough), they cancel each other out. This results in a region of zero electron density (a nodal plane) between the nuclei. This type of interaction forms an antibonding molecular orbital (ABMO). Electrons in ABMOs destabilize the molecule and have higher energy than the original atomic orbitals.

For two atomic orbitals to combine effectively and form molecular orbitals, three crucial conditions must be met:

Comparable Energy: — The atomic orbitals must have similar energy levels. For example, a 1s orbital from one atom will combine effectively with a 1s orbital from another atom, but not with a 2p orbital, as their energy difference is too large.
Proper Symmetry: — The atomic orbitals must have the correct orientation and symmetry to overlap effectively. A head-on overlap forms a sigma ($sigma$) bond, while a sideways overlap forms a pi ($pi$) bond. If the symmetries don't match, there's no net overlap, and no molecular orbital formation.
Maximum Overlap: — The atomic orbitals must overlap to a significant extent. The greater the overlap, the stronger the bond and the more stable the resulting molecular orbital.

Once formed, these molecular orbitals are filled with electrons according to the same rules that apply to atomic orbitals: the Aufbau principle (fill lowest energy first), Pauli's exclusion principle (max two electrons per orbital with opposite spins), and Hund's rule (maximize spin multiplicity for degenerate orbitals). The way these electrons fill the MOs determines the molecule's stability, bond order, and magnetic properties.