4.3: Double and Triple Bonds

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So far, we have considered what are known as single bonds; that is, all the C–C and C–H bonds in alkanes, and all the bonds in diamond. Each single bond involves two (and only two) electrons that are described by a bonding molecular orbital. In such a bonding orbital, most of the electron density is located between the two bonded atoms in a linear sigma (σ) bond. We have, however, already discussed albeit briefly bonds that involve more than one pair of electrons, namely those found in graphite. Recall that for graphite and graphene the bonds between carbon atoms in the sheet plane involve hybridized orbitals that are mixtures of the 2s² and 2p_x and 2p_y (that is sp² hybrid orbitals) leaving an unhybridized 2p_z orbital. On bonding, these unhybridized 2p_z orbitals reorganize to form what is known as a pi (π) bonding orbital. In π orbitals, the electron density lies above and below the axis connecting the bonded atoms. The combination of σ and π bonding orbitals produces a double bond. Double bonds are indicated by two lines, for example as in CH₂=CH₂(ethene).