Linus Pauling · J. Am. Chem. Soc. 53(4), 1367–1400 · April 1931 · California Institute of Technology
The problem
G. N. Lewis (1916) had pictured the chemical bond as a pair of electrons shared between two atoms, but the picture could not say why a shared pair binds, why bonds adopt definite angles, or why a saturated carbon atom is tetrahedral. Heitler and London (1927) had just answered the first question for the hydrogen molecule using the new quantum mechanics. Pauling set out to carry that result into the whole of structural chemistry.
Six rules for the electron-pair bond
The paper opens by laying down six rules. The first three restate the shared-pair idea in quantum-mechanical terms: a bond forms from one unpaired electron on each of two atoms; the two electron spins pair and cancel; and the paired electrons then take no further part in bonding. The last three are new and quantitative — the bond's energy comes from the resonance (exchange) of the two electrons between the atoms; of the orbitals available, those of lowest energy form the strongest bonds; and, decisively, a bond tends to lie in the direction in which the atom's bonding orbital is most concentrated, because that direction maximises the overlap of the two orbitals and so the strength of the bond.
Hybrid bond orbitals and the tetrahedral carbon
From the maximum-overlap rule Pauling drew his most consequential result. An atom need not bond through its bare s and p orbitals; it can mix them into equivalent hybrid orbitals that are better directed and overlap more. One s orbital combined with carbon's three 2p orbitals yields four equivalent hybrids pointing to the corners of a regular tetrahedron, mutually 109.47° apart — the long-mysterious tetrahedral carbon of organic chemistry, now derived from quantum mechanics. Mixing s with two p orbitals gives three coplanar bonds at 120°; with one p orbital, two opposed bonds at 180°.
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A magnetic criterion
The second half of the paper uses paramagnetic susceptibility — a magnetic measurement of the number of unpaired electrons — as an experimental criterion for the type of bonding in complexes of the transition metals, distinguishing essentially ionic from covalent (electron-pair) bonds. This is the origin of the 'inner-orbital' versus 'outer-orbital' distinction still used for coordination compounds.
What followed
This was the first of a celebrated series. The electronegativity scale and the concept of resonance came in the papers that followed and in the 1939 book that grew out of them. The directed-valence idea was reached independently the same year by the physicist J. C. Slater.
California Institute of Technology · 1931