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Ch400Ch7LN4 - Valence Bond Theory and Molecular Orbital...

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Valence Bond Theory and Molecular Orbital Theory Covalent Bonding and Orbital Overlap: Valence Bond Theory Lewis structures and VSEPR theory give us the shape of the molecule and the location of electrons in a molecule. They do not explain why a chemical bond forms. How do we account then, for molecular shape in terms of quantum mechanics? That is, which orbitals are involved in bonding? We use valence-bond theory : A covalent bond forms when the orbitals on two atoms overlap . The shared region of space between the orbitals is called the orbital overlap . There are two electrons (usually one from each atom) of opposite spin in the orbital overlap. As two nuclei approach each other their atomic orbitals (AO’s) overlap. The more the orbitals overlap, the stronger the bond. When p or d orbitals are involved in bonding, the resulting covalent bond has direction. Hybrid Orbitals The above works well for some simple molecules but fails miserably for many others. For example, look at CH 4 . Looking at the orbital energy diagram for C, what do you see? The valence 2s electrons fill an orbital, while the 2 2p electrons occupy different orbitals. This means that there are only 2 unpaired electrons, so you would think that C could only form 2 bonds. But C forms 4 bonds! How does this happen? To explain this, chemists developed a theory to explain how valence orbitals interact to form bonds. sp Hybrid Orbitals Consider the BeF 2 molecule: Be has a 1 s 2 2 s 2 electron configuration. There are no unpaired electrons available for bonding. So a simple AO overlap is not adequate to describe this molecule.
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