How did quantum mechanics improve explanations of chemical bonding after Gilbert N. Lewis?

While Gilbert N. Lewis’s 1916 model of the electron-pair bond was foundational, it was essentially static and lacked a quantitative basis to explain bond strengths, lengths, or molecular shapes. Quantum mechanics transformed these concepts by providing a rigorous physical framework for chemical bonding.123

Key advancements included:

• The Valence Bond (VB) Method: Developed by Linus Pauling and John Slater, this approach used quantum mechanical exchange effects to explain how electron spins allow atoms to bond. It introduced the concept of hybridization, which accounted for the specific spatial directionality of bonds—such as the tetrahedral structure of carbon—that Lewis’s model could not predict.34
• Molecular Orbital (MO) Theory: Pioneered by Robert Mulliken and others, this theory treats electrons as occupying wave functions that are spread over the entire molecule rather than being localized between two atoms. This approach proved more amenable to computational implementation and is now the standard for investigating molecular structures.53
• Addressing Exceptions: Quantum mechanics explained phenomena that defied the octet rule, such as electron-deficient compounds (e.g., diborane), by demonstrating that electron pairs can bond more than two atoms simultaneously.2