## Lecture 10/23 on Hybridization of Ethene

$sp, sp^{2}, sp^{3}, dsp^{3}, d^{2}sp^{3}$

jgreynoso 2J
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Joined: Fri Sep 25, 2015 3:00 am

### Lecture 10/23 on Hybridization of Ethene

Mr. Lavelle showed the lewis structure of both acethylene and ethene today. I am a little confused about the delocalized pi bonds. For acethylene with an sp hybridization, Lavelle illustrated that it contains two electrons occupying different orbitals with parallel spin in the 2p. Whereas for ethene, it has an sp2 hybridization where there is one electron occupying one orbital in the 2p. His diagrams illustrated the formation of a pi bond with both, but what does the difference between the delocalized pi bonds create? Is there multiple for the acethylene the diagram left out?

Thank you

Shaye Busse 3B
Posts: 20
Joined: Fri Sep 25, 2015 3:00 am

### Re: Lecture 10/23 on Hybridization of Ethene

For acetylene, or rather ethyne, there are two regions of electron density surrounding each carbon atom, meaning they will have 2sp hybridization for its sigma-bonding orbitals. This leaves two electrons per carbon atom, each occupying an unpaired 2p orbital. This means on each carbon atom, there is one sigma bond formed from the end-to-end overlap between the two carbon atom and one sigma bond formed from the end-to-end overlap between the carbon and hydrogen atom. This leaves two unpaired 2p orbitals around each carbon atom. Because of this, the two unpaired 2p orbitals from each carbon atom will form two weaker pi bonds with each other, accounting for the triple bond on ethyne.

In ethene, there are three regions of electron density around each carbon atom rather than one. This means the sigma-bonding orbitals for the carbon atoms will have 2sp2 hybridization. After all the sigma bonding, there is one unpaired 2p orbital per carbon atom resulting in one pi bond between the carbon atoms. This is what accounts for the double bond in ethene.