# Alkenes: Properties

## Alkenes: Valence Bond Model

In principle, two models can be used to explain the electronic structure of the double bond.

1. Valence bond model (Generate hybrid orbitals, add one electron to each and combine to form bonds)
2. Molecular orbital theory (Combine atom or hybrid orbitals to form molecular orbitals and fill with electrons)
Fig.1
Construction principle for alkenes: $sp2$ Hybridization

In a simple and at least virtually more effective method, hybrid orbitals are formed and combined with unpaired of occupied to form bonds containing two paired electrons each. The four electrons in the valence shell of carbon are distributed over four atomic orbitals ( 2s, 2px, 2py, 2pz ). The equivalency of the four bonds in methane is explained by mixing all four atomic orbitals to form identical $sp3$-hybrid orbitals (s + 3 p = 4 $sp3$). In the case of ethene, one s orbital and two p orbitals of each C atom are combined to form three $sp2$ orbitals while the third p orbital remains unchanged. (s + 3 p = 3 $sp2$ + p).

An ethene molecule can now be constructed from two $sp2$-hybridized C atoms, each with three $sp2$ orbitals (gray), one p orbital (blue) and four s hydrogen orbitals.

One σ bond (gray) and one π bond (blue) form the bonding between both C atoms. σ Bonds are bonds in which the electron density is distributed symmetrically in respect to the rotation about the bond axis. This also includes the C-H bond, for example. π Bonds contain a so-called nodal plane which passes through both of the involved nuclei and in which the electron density equals zero. In this case, bonding is achieved by electrons that are located above and below the plane. Though this generates two electron paths (blue lobes in Fig.1) between the atoms, the π bond is counted as one bond.

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