Bonding

Organic Chemistry I
Bonding: Atomic Orbitals & Molecular Orbitals
Reference: Vollhardt & Schore, Organic Chemistry, W.H.Freeman
The electron is described by wave equations.
(A) A wave. The signs of the amplitude are assigned arbitrarily. At points of zero amplitude, called nodes, the wave
changes sign. (B) Waves with amplitudes of like sing (in phase) reinforce each other to make a larger wave. (C) Waves
out of phase subtract from each other to make a smaller wave.
Bonds are made by in-phase overlap of atomic orbitals.
In-phase (bonding) and out-of-phase (antibonding) combinations of 1s atomic orbitals. The + and - signs denote the sign
of the wave function, not charges. Electrons in bonding molecular orbitals have a high probability of occupying the space
between the atomic nuclei, as required for good bonding. The antibonding molecular orbital has a nodal plane where the
probability of finding electrons is zero. Electrons in antibonding molecular orbitals are most likely to be found outside the
space between the nuclei and therefore do not contribute to bonding.
!2
Schematic representation of the interaction of (A) two singly occupied atomic orbitals (as in H2) and (B) two doubly
occupied atomic orbitals (as in He2) to give two molecular orbitals (MO). (These diagrams are not drawn to scale.)
Formation of an H-H bond is favorable because it stabilizes two electrons. Formation of an He-He bond stabilizes two
electrons (in the bonding MO) but destabilizes two others (in the antibonding MO). Bonding between He and He thus
results in no net stabilization. Therefore, He is monoatomic.
The overlap of atomic orbitals gives rise to sigma and pi bonds.
Bonding between atomic orbitals. (A) 1s and 1s (e.g., H2), (B) 1s and 2p (e.g., HF), (C) 2p and 2p (e.g., F2), (D) 2p and
3p (e.g., FCl) aligned along internuclear axes, σ bonds; (E) 2p and 2p perpendicular to internuclear axis (e.g., H2C=CH2),
a π bond. Note the arbitrary use of + and - signs to indicate in-phase interactions of the wave functions.