I the electron density of a true π bond is symmetric

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I The electron density of a true π bond is symmetric about the nodal plane. I This requires that the nodal plane coincide with the molecular plane. Any out-of-plane bending (e.g. in fullerenes) will break the symmetry required to make a true π bond. I In molecules where we can’t form a true π bond, σ - π separation is only an approximation, but usually a good one. I π bonds are often important for understanding reactivity in organic chemistry, so it’s convenient to be able to work with them without worrying about the σ bonding.
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Ethyne C C H H I Linear geometry around each carbon = sp hybridization I Each carbon atom has two p orbitals left over: H H C C I These p orbitals combine into two π bonds.
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Formaldehyde C H H :O: I Trigonal planar geometry at the carbon atom = sp 2 hybridization I Two views of how to treat the O atom: I The O atom has a trigonal planar electronic geometry sp 2 hybridized. I The O atom is terminal hybridization is unnecessary. The O atom can form a σ bond using a p orbital.
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C H H :O: sp 2 -sp 2 or 2p-sp 2 valence bond 2 1s-sp valence bonds I Either way, the carbon and O atoms each have one p orbital left over which can used to form the π bond.
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