Chemistry · Ch 10 — Chemical Bonding
Molecular Orbital Theory
Molecular Orbital Theory
Both Lewis theory and valence bond theory give useful, largely correct qualitative pictures of chemical bonding and molecular structure -- but both are built on the idea of electrons pairing up into BONDS, and this makes both theories inadequate for certain observed molecular properties. The textbook example: both theories predict that the oxygen molecule should be DIAMAGNETIC (all electrons paired, as in the σ+π picture of section 10.8.2). But experimentally, LIQUID oxygen is observed to be attracted toward the poles of a strong magnet -- meaning oxygen is actually paramagnetic (it has unpaired electrons). Because Lewis theory and VB theory both treat bond formation strictly in terms of electron PAIRS, neither can explain this paramagnetic behaviour.
F. Hund and Robert S. Mulliken developed a different bonding theory -- molecular orbital theory (MOT) -- that DOES explain the magnetic behaviour of molecules such as O₂. Its salient features:
- When atoms combine to form a molecule, their INDIVIDUAL atomic orbitals lose their separate identity and merge into entirely NEW orbitals, called molecular orbitals.
- The shapes of the resulting molecular orbitals depend on the shapes of the atomic orbitals that combined to form them.
- The NUMBER of molecular orbitals formed equals the number of combining atomic orbitals. HALF of the resulting molecular orbitals have LOWER energy than the original atomic orbitals -- these are called bonding molecular orbitals -- while the other half have HIGHER energy -- these are called antibonding molecular orbitals. Bonding MOs are denoted σ (sigma), π (pi), δ (delta); the corresponding antibonding MOs are denoted σ*, π*, δ* (with an asterisk).
- Electrons in a molecule are filled into these newly-formed molecular orbitals following the SAME rules used for filling atomic orbitals: the Aufbau principle (lowest energy orbital fills first), Pauli's exclusion principle (no two electrons in the same orbital with the same spin), and Hund's rule (degenerate orbitals are singly occupied first, with parallel spins, before any is doubly occupied). …
Bonding in Some Heteronuclear Diatomic Molecules
The same LCAO/Aufbau/Pauli/Hund bookkeeping applies to HETEROnuclear diatomics too, though now the two combining atoms are different, so their atomic-orbital energy levels are not identical (each side of the energy-level diagram is drawn from a different element's orbitals).
Carbon monoxide, CO (Fig 10.36(a)). Atomic configurations: C, 1s² 2s² 2p²; O, 1s² 2s² 2p⁴. CO has 14 total electrons (6 + 8), exactly the same electron count as N₂ -- CO is said to be isoelectronic with N₂, and its molecular-orbital configuration mirrors N₂'s accordingly: . Bond order -- a triple bond, matching CO's Lewis structure (). All electrons paired → diamagnetic.
Nitric oxide, NO (Fig 10.36(b)). Atomic configurations: N, 1s² 2s² 2p³; O, 1s² 2s² 2p⁴. NO has 15 total electrons -- ONE more than CO/N₂ -- and this extra electron must go into the next available (antibonding) orbital. Molecular-orbital configuration of NO: . Bond order -- a FRACTIONAL bond order, something neither Lewis theory nor simple bond-counting could ever produce, but which molecular orbital theory handles naturally. The one unpaired electron (in π*2p_y) makes NO paramagnetic. …
What this figure shows. Energy-level diagram for CO, built from carbon's and oxygen's atomic orbitals (labelled separately on each side since the two atoms differ): σ2s, σ2s filled, the degenerate π2py/π2pz level fully filled, and σ2px filled -- the same bonding-MO occupation pattern as N₂, all antibonding π/σ …
What this figure shows. Energy-level diagram for NO, built from nitrogen's and oxygen's atomic orbitals: the same filled pattern as CO/N₂ (σ2s, σ2s, π2py, π2pz, σ2px) PLUS one additional electron sitting alone in the antibonding π2py orbital -- the odd, unpaired electron responsible for NO's p …