MO theory predicts the distribution of electrons in a molecule. This helps predict properties such as bond order, shape, and electrical properties.
Valence Bond (VB) theory and Lewis Structures explain the properties of simple molecules. Molecular Orbital (MO) theory better explains the properties of more complex molecules.
MO theory explains the partial bonds of NO₃⁻ without using resonance. This makes MO theory more useful for the description of extended π systems.
Also, in benzene the six p electrons are in three molecular π orbitals around the ring. These orbitals spread the six electrons evenly over all the carbon atoms. All carbon-carbon bonds in benzene are equivalent. There is no need to use resonance.
MO theory also explains other properties better than VB theory.
Whenever you can draw two or more Lewis structures for a molecule, the actual structure is none of the structures but is a resonance hybrid of them all. For example, we say that the two resonance structures of ozone contribute to a resonance hybrid in which the π electrons are delocalized between all three O atoms. This makes the two O-O bonds half way between single and double bonds or 1.5 bonds.
Molecular Orbital theory starts by assuming that the three atomic p orbitals on the O atoms overlap to form three molecular π orbitals that extend over the whole molecule. We end up with two electrons in a bonding π orbital; two electrons in a nonbonding #π^n# orbital; and no electrons in an antibonding #π^✳# orbital.
Calculations show that the #π# bond order of each O-O bond is 0.5. When we add the underlying σ bond, the total O-O bond order is 1 + 0.5 = 1.5. Thus, Molecular Orbital theory explains resonance delocalization automatically as the natural state of the molecule.