Problem 58 - Chemistry: The Central Science

Question

(a) Write a single Lewis structure for \(\mathrm{SO}_{3}\), and determine the hybridization at the \(\mathrm{S}\) atom. (b) Are there other equivalent Lewis structures for the molecule? (c) Would you expect \(\mathrm{SO}_{3}\) to exhibit delocalized \(\pi\) bonding? Explain.

Step-by-Step Solution

Answer

The Lewis structure for SO3 has a central Sulfur atom with three double bonds to three Oxygen atoms, and it exhibits sp2 hybridization. There are two more equivalent resonance structures possible for SO3, suggesting delocalized π bonding as the π electrons are spread across the entire molecule.

Step 1: Write a Lewis structure for SO3

To construct the Lewis structure for SO3, we will first find out the total number of valence electrons in the molecule: Valence electrons for Sulfur: 6 (as it belongs to group 6 in the periodic table) Valence electrons for each Oxygen: 6 (as it belongs to group 6 in the periodic table) Total valence electrons = 1(S) * 6(valence electrons for S) + 3(O) * 6(valence electrons for each O) = 6 + 18 = 24 Now, arrange the atoms and distribute the electrons: 1. Place the least electronegative atom (Sulfur) in the center and connect it with three Oxygen atoms using single bonds. 2. Distribute the remaining valence electrons to complete the octets for the outer atoms. 3. If the central atom (Sulfur) does not have an octet, use lone pair electrons from the neighboring atoms to create double or triple bonds.

Step 2: Determine the hybridization of the Sulfur atom

To determine the hybridization of an atom, we need to know the number of sigma bonds and lone pairs around that atom. In SO3, the Sulfur atom is connected to each of the three Oxygen atoms with a double bond. Therefore, it has no lone pairs and three sigma bonds. Hybridization can be found using the following formula: hybridization = number of sigma bonds + lone pairs hybridization = 3 sigma bonds + 0 lone pairs = 3 The sum is equal to 3, which means that the hybridization of the Sulfur atom in SO3 is sp2.

Step 3: Check for other equivalent Lewis structures

To determine if there are other equivalent Lewis structures for SO3, we can try drawing other resonance structures with different arrangements of double bonds. Upon trying, we can find that there are indeed two more equivalent resonance structures possible for SO3. The three structures are different only in the arrangement of double bonds between the Sulfur and Oxygen atoms. Thus, they are all equivalent Lewis structures for the molecule.

Step 4: Analyze the possibility of delocalized π bonding

Delocalized π bonding happens when π electrons are spread across more than two atoms in a molecule. In the case of SO3, we have multiple resonance structures. This indicates that the π electrons are spread across the entire molecule. Since the double bonds are not confined to one location but are distributed between different Oxygen and Sulfur atoms, we can conclude that SO3 exhibits delocalized π bonding.