Tetravalence and shapes of carbon compounds

Carbon always forms four bonds (it is tetravalent). How it uses its orbitals depends on what kind of bonds it makes:

• Four σ bonds (only single bonds): sp³, tetrahedral, bond angle 109.5°, as in methane, CH₄.
• Three σ + one π (one double bond): sp², trigonal planar, 120°, as in ethene, C₂H₄.
• Two σ + two π (a triple bond, or two double bonds): sp, linear, 180°, as in ethyne, C₂H₂.

A quick way: count the σ bonds on the carbon: 4 means sp³, 3 means sp², 2 means sp. In propene, CH₃–CH=CH₂, the CH₃ carbon is sp³, and the two carbons of the double bond are sp².

The s character of a hybrid is the share of s orbital in it: sp 50%, sp² 33%, sp³ 25%. An s orbital holds electrons closer to the nucleus than a p orbital, so more s character means shorter, stronger bonds, and a carbon that is more electronegative.

A π bond forms by sideways overlap of parallel p orbitals. To keep them parallel, all six atoms of ethene, H₂C=CH₂, must lie in one plane, and the two ends cannot rotate freely about the C=C bond without breaking the π bond.

The π electron cloud lies above and below the plane of the molecule, out in the open, where attacking reagents can easily reach it. So π bonds are the reactive centres of many organic molecules.

Remember: each single bond is one σ; a double bond is σ + π; a triple bond is σ + 2π.