Molecular Orbital Theory

The trap that costs marks: students fill molecular orbital diagrams correctly but then count only bonding electrons when calculating bond order — forgetting that antibonding electrons subtract. For F₂, this gives bond order 4 instead of the correct 1. A second common confusion is the MO energy ordering: for B₂, C₂, and N₂, the σ2p_z orbital sits above the π2p orbitals (due to s–p mixing), but for O₂, F₂, and Ne₂ the σ2p_z drops below the π2p orbitals. Using the wrong ordering leads to incorrect electron configurations, wrong bond orders, and wrong magnetic-property predictions.

What MOT actually says. When two atomic orbitals combine (LCAO — Linear Combination of Atomic Orbitals), they produce two molecular orbitals: one bonding (lower energy, constructive overlap) and one antibonding (higher energy, destructive overlap, marked with an asterisk *). Electrons fill MOs following Aufbau, Pauli exclusion, and Hund's rule — exactly as in atoms. Bond order is then calculated as:

BO = (N_b − N_a) / 2

where N_b = total bonding electrons and N_a = total antibonding electrons (NCERT Class 11 Chemistry Chapter 4, page 24).

The ordering switch you must memorise. For homonuclear diatomics up to and including N₂ (Z ≤ 7): σ1s < σ1s < σ2s < σ2s < π2p_x = π2p_y < σ2p_z < π2p_x = π2p_y < σ2p_z. From O₂ onward (Z ≥ 8): σ2p_z drops below the π2p pair. The boundary is at the N₂/O₂ transition. Getting this wrong flips paramagnetism predictions — O₂ is paramagnetic (two unpaired electrons in π), and this is a frequent NEET question.

Watch out: if BO = 0, the molecule does not exist (e.g., He₂ with BO = 0; He₂⁺ with BO = 0.5 does exist). A fractional bond order is valid and physically meaningful.

---