Describe the bonding in \(\mathrm{SO}_{2}\) and \(\mathrm{SO}_{3}\) using the localized electron model (hybrid orbital theory). How would the molecular orbital model describe the \(\pi\) bonding in these two compounds?

In both SO₂ and SO₃, sulfur undergoes sp² hybridization, forming sigma bonds with each oxygen atom using its sp² hybrid orbitals. In SO₂, there's a lone pair on sulfur, and a double bond is formed between the sulfur and each oxygen atom due to two π bonds. In SO₃, a trigonal planar geometry is formed with three equivalent S-O resonance structures due to the delocalization of electrons in π bonds. According to the molecular orbital model, the π bonding in SO₂ and SO₃ involves the combination of p-orbitals on sulfur and oxygen atoms, creating molecular orbitals. In SO₂, one bonding and one antibonding π molecular orbitals are formed. In SO₃, a set of four π molecular orbitals is formed, with three bonding and one antibonding orbitals, resulting in delocalized π bond formation and resonance.