Methane

Methane is the simplest organic molecule: one carbon holding four hydrogens. It’s the starting point for almost everything else, because the way its carbon bonds is the way carbon bonds in every saturated chain and ring you’ll meet later.

Carbon has four valence electrons: two in its 2s orbital and two in 2p orbitals. To make four bonds it uses all four of those orbitals, the one 2s and the three 2p. If it used them as they are, you’d expect one bond (the one from the 2s orbital) to be different from the other three. But all four C–H bonds in methane are identical. The explanation is hybridization: the 2s and the three 2p orbitals mix into four equivalent sp³ hybrid orbitals, each one part s and three parts p. Each sp³ orbital overlaps head-on with a hydrogen 1s orbital to make a σ (sigma) bond.

Four electron pairs push each other as far apart as they can get, so the four bonds point to the corners of a tetrahedron. Every H–C–H angle is about 109.5°. Rotate the model and look down any one C–H bond: the other three hydrogens form a perfect three-legged stand. Turn on Angles & lengths to see the measured angle.

Methane has no π bonds and no leftover p orbitals, so there is nothing to show for the p-orbital and π-electron layers. Its electron density is a smooth, nearly spherical cloud, which is one reason methane molecules barely stick to each other and methane is a gas at room temperature.