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Benzene, and aromatic compounds in general, is interesting because its properties are quite different from conjugated (double-single alternating bonds) alkene compounds.
In the case of benzene, the Lewis Bonding model is not very useful in predicting properties because there is a significant effect from the behavior of non-hybridized electron-orbitals, which is outside the scope of the Lewis model. Instead, we look to the Atomic Hybrid Orbital Model for a clue regarding the unusual properties of benzene. That clue lies in the behavior of unhybridized p-orbitals and their spatial relationship to sp2 hybrid orbitals. The geometry of sp2 hybrid orbitals are planar (flat) while the unhybridized p-orbital is orthogonal (a straight up and straight down arrangement relative to the molecular plane). Within the continuous-conjugated benzene ring, p-orbitals of separate carbon nucleii are close enough to overlap and, because of their orthogonal arrangement, the overlap is side-to-side: Pi-Bonds are formed. Benzene consists of a continuous-array of six overlapping p-orbitals, which behave like a single molecular orbital (yet another theory of bonding in molecules), which provides a stabilizing effect within the benzene molecule.
Normally, double bonds (such as in propylene) can be easily hydrogenated (addition of hydrogen), but this is not the case with benzene. This behavior of benzene (not allowing addition of atoms across its double bonds) is the property that the Lewis Model cannot predict. The hybrid orbital model predicts the property, just discussed, on the grounds of energetic unfavorability, with regard to the hypothetical resulting geometry of added atoms or groups (such as the hydroxyl group). This hypothetical resulting geometry is a tight-angle arrangement (less than 104 deg.) at the sites of atom addition across the affected carbon atoms. The following illustration provides a summary.
Thank you for reading!
Benzene, and aromatic compounds in general, is interesting because its properties are quite different from conjugated (double-single alternating bonds) alkene compounds.
In the case of benzene, the Lewis Bonding model is not very useful in predicting properties because there is a significant effect from the behavior of non-hybridized electron-orbitals, which is outside the scope of the Lewis model. Instead, we look to the Atomic Hybrid Orbital Model for a clue regarding the unusual properties of benzene. That clue lies in the behavior of unhybridized p-orbitals and their spatial relationship to sp2 hybrid orbitals. The geometry of sp2 hybrid orbitals are planar (flat) while the unhybridized p-orbital is orthogonal (a straight up and straight down arrangement relative to the molecular plane). Within the continuous-conjugated benzene ring, p-orbitals of separate carbon nucleii are close enough to overlap and, because of their orthogonal arrangement, the overlap is side-to-side: Pi-Bonds are formed. Benzene consists of a continuous-array of six overlapping p-orbitals, which behave like a single molecular orbital (yet another theory of bonding in molecules), which provides a stabilizing effect within the benzene molecule.
Normally, double bonds (such as in propylene) can be easily hydrogenated (addition of hydrogen), but this is not the case with benzene. This behavior of benzene (not allowing addition of atoms across its double bonds) is the property that the Lewis Model cannot predict. The hybrid orbital model predicts the property, just discussed, on the grounds of energetic unfavorability, with regard to the hypothetical resulting geometry of added atoms or groups (such as the hydroxyl group). This hypothetical resulting geometry is a tight-angle arrangement (less than 104 deg.) at the sites of atom addition across the affected carbon atoms. The following illustration provides a summary.
Thank you for reading!
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