Ethene

In ethene (also called ethylene) each carbon bonds to only three atoms: two hydrogens and the other carbon. So each carbon mixes its 2s orbital with just two of its 2p orbitals, making three sp² hybrid orbitals. Those three lie flat in one plane, about 120° apart. They make the C–H bonds and a C–C σ bond.

That leaves one 2p orbital on each carbon, standing straight up and down from the plane. Turn on p orbitals and look edge-on: the two p orbitals sit side by side, parallel, and they overlap sideways. That sideways overlap is the π bond. A double bond is one σ bond plus one π bond.

Now turn on π electrons. The π bond’s electrons sit in two clouds, one above the plane of the atoms and one below, not along the line between the carbons. The faint sheet marks that plane: the π density is exactly zero there. This is the key picture for alkenes: the π electrons are exposed above and below the molecule, which is why alkenes react with electron-seeking reagents.

The π bond also explains two other facts. Ethene is completely flat, because the p orbitals can only overlap if both CH₂ ends lie in the same plane. And the two ends can’t rotate without breaking that overlap, which is why cis and trans isomers of alkenes exist. The C=C bond (about 1.33 Å) is also shorter than ethane’s C–C bond.