Greetings Valued Readers,
This post will focus on the Hybrid Orbital Bonding Theory as it applies to Benzene. Benzene owes its great stability to its conjugated bond system, which provides continuous orbital overlap all the way around the carbon ring.
Molecular Orbital Theory provides the idea that a continuous orbital spanning the full extent of a molecular ring may exist. The Hybrid Orbital Model provides a similar explanation, but with the idea that individual atomic orbitals will "mix" to form hybrid atomic orbitals as covalent bonding occurs. The idea is that the formation of hybrid orbitals happens so that the bonding orbitals will have a geometric arrangement matching that of the molecule itself.
If the geometry of bonding orbitals did not closely match the molecule geometry, then the bonding would be strained. This strained covalent bonding would consist of electrons in a higher potential energy state, which in-turn would cause the molecule to be less stable.
The type of hybrid orbital bonding for the benzene ring system is sp2. The sp2 mixed orbitals occur at each carbon atom. The number of hybrid orbitals formed is always equal to the number of individual atomic orbitals prior to mixing. So, a combination of one s-orbital and two p-orbitals means that three sp2 hybrid orbitals will form. According to Valence Shell Electron-Pair Repulsion (VSEPR) theory, three hybrid orbitals will have a trigonal planar geometry, with a bond angle of 120 degrees. This is exactly the geometry required for the benzene ring!
There is yet more to this story, however. There is one p-atomic orbital left after the three sp2 orbitals are formed at each carbon atom. VSEPR theory, once again, tells us that this remaining p-orbital must be arranged perpendicular to the plane of the benzene ring for greatest stability. That is exactly what happens and the lone p-atomic orbital at each carbon atom also participates in carbon to carbon bonding.
Two of the three sp2 orbitals at each carbon atom participate in bonding within the ring system. The geometry of these orbitals allows for an end-to-end overlap resulting in the formation of sigma bonds. The lone p-atomic orbitals also bond, but because they are all perpendicular to the plane of the ring, they bond side-to-side (forming pi bonds).
The following graphic illustrates the bonding for benzene.
The diagram also shows orbital overlap for each hydrogen atom of the benzene ring. That bonding is also of the sigma type because it involves a direct overlap of the hydrogen s-orbital with a sp2 hybrid orbital extending from the benzene ring.
That's all for now and, as always, thank you for reading!
A Publication of http://ExcellenceInLearning.biz
This post will focus on the Hybrid Orbital Bonding Theory as it applies to Benzene. Benzene owes its great stability to its conjugated bond system, which provides continuous orbital overlap all the way around the carbon ring.
Molecular Orbital Theory provides the idea that a continuous orbital spanning the full extent of a molecular ring may exist. The Hybrid Orbital Model provides a similar explanation, but with the idea that individual atomic orbitals will "mix" to form hybrid atomic orbitals as covalent bonding occurs. The idea is that the formation of hybrid orbitals happens so that the bonding orbitals will have a geometric arrangement matching that of the molecule itself.
If the geometry of bonding orbitals did not closely match the molecule geometry, then the bonding would be strained. This strained covalent bonding would consist of electrons in a higher potential energy state, which in-turn would cause the molecule to be less stable.
The type of hybrid orbital bonding for the benzene ring system is sp2. The sp2 mixed orbitals occur at each carbon atom. The number of hybrid orbitals formed is always equal to the number of individual atomic orbitals prior to mixing. So, a combination of one s-orbital and two p-orbitals means that three sp2 hybrid orbitals will form. According to Valence Shell Electron-Pair Repulsion (VSEPR) theory, three hybrid orbitals will have a trigonal planar geometry, with a bond angle of 120 degrees. This is exactly the geometry required for the benzene ring!
There is yet more to this story, however. There is one p-atomic orbital left after the three sp2 orbitals are formed at each carbon atom. VSEPR theory, once again, tells us that this remaining p-orbital must be arranged perpendicular to the plane of the benzene ring for greatest stability. That is exactly what happens and the lone p-atomic orbital at each carbon atom also participates in carbon to carbon bonding.
Two of the three sp2 orbitals at each carbon atom participate in bonding within the ring system. The geometry of these orbitals allows for an end-to-end overlap resulting in the formation of sigma bonds. The lone p-atomic orbitals also bond, but because they are all perpendicular to the plane of the ring, they bond side-to-side (forming pi bonds).
The following graphic illustrates the bonding for benzene.
The diagram also shows orbital overlap for each hydrogen atom of the benzene ring. That bonding is also of the sigma type because it involves a direct overlap of the hydrogen s-orbital with a sp2 hybrid orbital extending from the benzene ring.
That's all for now and, as always, thank you for reading!
A Publication of http://ExcellenceInLearning.biz
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