For a molecule, we can calculate the bond order to characterize the bonding in the molecule. The bond order is equal to half of the sum of the number of electrons in bonding orbitals minus the number of electrons in anti-bonding orbitals. For diatomics, the bond order will tell us if we have single, double, or triple bonds. For example, below is, again, our MO diagram for hydrogen.
Hydrogen has two electrons in a bonding orbital and zero electrons in anti-bonding orbitals. Therefore the bond order is one. [ B.O = 0.5(2-0) = 1].
For larger molecules, the idea is more difficult to translate as MO deals with the whole molecule rather than individual bonds. For example, below is a picture of the molecular orbitals of methane.
You can see the atomic orbitals from carbon (2s and 2p) and four 1s orbital from hydrogen. The MOs are in the middle. Methane has four σ MOs and four σ* MOs. All of the electrons are in the bonding orbitals. The bond order in methane is then = 0.5(8-0) = 4. This is not because there is a quadruple bond. Instead, this is because MO theory looks at the whole molecule. We would describe methane as having four C-H single bonds. Thus the bond order is four. However, you have to remember that MO looks at the whole molecule not local bonds. This is closer to "reality" as far as QM, but it is confusing conceptually. That is why as chemists we use MO to find geometries and energies, but we talk about bonds and local bonding theories.