Complete the following resonance structures for \(\mathrm{POCl}_{3}\) a. Would you predict the same molecular structure from each resonance structure? b. What is the hybridization of P in each structure? c. What orbitals can the P atom use to form the \(\pi\) bond in structure \(\mathrm{B}\) ? d. Which resonance structure would be favored on the basis of formal charges?

To draw the Lewis structures, first, count the total number of valence electrons. P has 5 valence electrons, O has 6, and each Cl has 7. This sums to 5+(6+7×3)=32 electrons required for all bonds and lone pairs. Draw a skeleton structure (with single bonds) and distribute electrons as needed to fulfill the octet rule. Draw alternative structures by moving electrons between atoms to make different types of bonds between the same atoms. Here are the resonance structures for \(\mathrm{POCl}_{3}\): Structure A: O ‖ Cl — P — Cl | Cl Structure B: O \ Cl — P — Cl | Cl

a. Would you predict the same molecular structure from each resonance structure? Yes, the molecular structure resulting from each resonance structure would be the same. Each resonance structure has the same arrangement of atoms and different sharing of electrons, resulting in a distorted tetrahedral structure. Because resonance structures are averaged, the electrons are distributed between the different bonding arrangements, which would not change the geometry.

b. What is the hybridization of P in each structure? Recall that the hybridization of an atom corresponds to the number and types of atomic orbitals forming hybrid orbitals. To determine the hybridization of P, we need to analyze the types of electron pairs around the P atom. In both Structure A and B, P is surrounded by 4 bonding pairs and no lone pairs, resulting in a steric number of 4. Therefore, P must use one s orbital and three p orbitals (sp3 hybridization) in both structures.

c. What orbitals can the P atom use to form the \(\pi\) bond in structure B? In structure B, there is a \(\pi\) bond between P and O. The \(\pi\) bond is formed from the electron pairs not involved in the hybrid orbitals (e.g., electrons from the remaining unhybridized p orbitals). In this case, the P atom forms a \(\pi\) bond with the O atom using its unhybridized 3p orbital to interact with the p orbital of the O atom.

d. Which resonance structure would be favored on the basis of formal charges? To determine the favored resonance structure, we need to calculate the formal charge for each atom in the structures. The formal charge is calculated using this formula: Formal charge = (number of valence electrons) - (number of lone pair electrons) - 0.5 × (number of bonded electrons) Calculate the formal charges for each atom in both structures: Structure A: P: 5 - 0 - 0.5 × (8) = +1 O: 6 - 4 - 0.5 × (4) = 0 Cl: (each) 7 - 6 - 0.5 × (2) = 0 Structure B: P: 5 - 0 - 0.5 × (10) = 0 O: 6 - 4 - 0.5 × (6) = 0 Cl: (each) 7 - 6 - 0.5 × (2) = 0 For the most stable structure, we want the formal charges on the atoms to be as close to zero as possible. Based on the formal charges calculated, Structure B has a formal charge of 0 for all atoms, which is closer to zero than in Structure A. Therefore, Structure B would be the favored resonance structure on the basis of formal charges.