Chemical Bonding

Two orbitals can meet in two ways. A sigma bond is head-on overlap along the internuclear axis, so the density sits on that axis and the atoms can spin freely around it. A pi bond is sideways overlap of two parallel $p$ orbitals, with lobes meeting above and below the nuclear plane and a nodal plane carrying nothing on the axis. That lateral overlap is weaker, so a pi bond never travels alone: a single bond is one sigma, a double bond one sigma plus one pi, a triple bond one sigma plus two pi.

Molecular orbital theory arrives at the same place from the other direction. Each pair of overlapping atomic orbitals throws up a bonding orbital of lower energy and an antibonding orbital of higher energy, so a linear combination of atomic orbitals gives both, and the count that falls out is bond order, $\frac{a-b}{2}$. Helium gives zero, which is why $\text{He}_2$ never forms. Nitrogen gives three, one sigma and two pi, and that triple bond is why nitrogen is far less reactive than oxygen. Oxygen gives two, and with two electrons unpaired in antibonding orbitals it is paramagnetic, which valence bond theory cannot account for.

In benzene the pi electrons spread above and below the ring rather than sitting between two atoms, so its molecular orbitals are delocalised and its sigma to pi ratio is $4:1$.