Orbital overlap: Difference between revisions

'''Orbital overlap''' is a concept used in theories of the [[chemical bond]]. It refers to the concentration of orbitals on adjacent atoms in the same region(s) of space, which can lead to bond formation. The importance of orbital overlap was emphasized by [[Linus Pauling]] to explain the molecular [[bond angle]]s observed through experimentation and is the basis for the concept of [[orbital hybridisation]]. ''s''&nbsp;orbitals are spherical and have no directionality while ''p'' orbitals are oriented 90° to one another. A theory was needed therefore to explain why molecules such as [[methane]] (CH₄) had observed bond angles of 109.5°.<ref>Anslyn, Eric V./Dougherty, Dennis A. (2006). ''Modern Physical Organic Chemistry''. University Science Books.</ref> Pauling proposed that s and p orbitals on the carbon atom can combine to form hybrids (sp³ in the case of methane) which are directed toward the hydrogen atoms. The carbon hybrid orbitals have greater overlap with the hydrogen orbitals, and can therefore form stronger C–H bonds.<ref>Pauling, Linus. (1960). ''The Nature Of The Chemical Bond''. Cornell University Press.</ref>

 

[[Category:Chemical bonding]]

[[Category:Molecular geometry]]

 

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