Draw reasonable resonance structures for the following ions: (a) \(\mathrm{HSO}_{4}^{-},\) (b) \(\mathrm{PO}_{4}^{3-},\) (c) \(\mathrm{HSO}_{3}^{-}\) (d) \(\mathrm{SO}_{3}^{2-}\). (Hint: See comment on Example \(\left.9.11 .\right)\)

Lewis structures are diagrams that show the distribution of valence electrons around atoms within a molecule. They serve as a visual representation of the molecular structure that helps in predicting the shape, reactivity, and properties of the molecule. A correct Lewis structure is crucial for understanding the molecule's bonding and its potential resonance structures. When drawing Lewis structures, one must account for electrons pairs surrounding the atoms, which can be in the form of bonding pairs (shared by two atoms) or lone pairs (belong to a single atom). Importantly, Lewis structures should adhere to the octet rule, meaning that most atoms strive to have eight electrons in their valence shell.

In the exercise for ions such as \(\mathrm{HSO}_4^-\), \(\mathrm{PO}_4^{3-}\), and \(\mathrm{HSO}_3^-\), we first sketch the Lewis structure considering these principles. This step sets a foundation for exploring possible resonance structures by showing us where the valence electrons are situated and how they are shared among the atoms.

Pi (π) electrons are located in pi bonds, which are formed by the sideways overlap of p orbitals on adjacent atoms. These electrons are often found in double and triple bonds and are highly mobile compared to electrons in single (sigma) bonds. Pi electrons can be delocalized across a molecule when there are resonance structures, making these molecules more stable due to the distribution of electron density.

In our exercise for polyatomic ions like \(\mathrm{HSO}_4^-\) and \(\mathrm{SO}_3^{2-}\), pi electrons are central to understanding their resonance structures. During resonance, pi electrons are not locked into one position but can be shuffled between different atoms, giving rise to equivalent structures that contribute to the overall description of the molecule.

Electron delocalization is a phenomenon where electrons, particularly pi electrons, are shared among multiple atoms in a molecule rather than being localized between a single pair of atoms. This delocalization occurs within a conjugated system or within ions where several resonance structures can be drawn.

Delocalization is significant because it increases the stabilization of the molecule or ion, lowering its potential energy. The classic example is benzene, where the pi electrons are not confined to a particular carbon-carbon bond but are spread evenly across the ring structure. In the context of the ions from the exercise—like the sulfate or sulfite ions—the possibility to draw multiple resonance structures illustrates this electron delocalization over the sulfur and oxygen atoms.

Stability and Geometry

The sulfate ion (\(\mathrm{SO}_4^{2-}\)) is a tetrahedral species with sulfate's central sulfur atom bonded to four oxygen atoms. The overall charge of -2 is delocalized across the four oxygen atoms. Despite having a formal charge, the sulfate ion is widely known for its stability, which arises from efficient electron delocalization.

Resonance Structures

When drawing resonance structures for the sulfate ion, multiple forms can be sketched, with the negative charges shifting between the oxygen atoms and double bonds being redistributed. These structures should depict all the possible ways that the pi electrons can be situated around the molecule, which are equally plausible and contribute to the overall description of the sulfate ion.

Structure Details

The phosphate ion (\(\mathrm{PO}_4^{3-}\)) is another example of a tetrahedral arrangement around the central phosphorus atom. This ion carries a -3 charge, which is shared among the phosphorus and oxygen atoms. Understanding its resonance structures is crucial for predicting interactions within biochemical systems, as phosphate groups are integral to energy-transfer molecules like ATP.

Multiple Forms

For the phosphate ion, the possible resonance structures involve shifting the pi electrons on the oxygen atoms and alternating double bonds with the central phosphorus atom. These various forms represent why the phosphate ion can be highly reactive and an essential component in biochemistry.

Understanding the Sulfite Ion

The sulfite ion (\(\mathrm{SO}_3^{2-}\)) is somewhat similar to the sulfate ion but with one less oxygen atom, resulting in a trigonal pyramidal shape. The charge delocalization and the possible resonance structures for the sulfite ion give clues about its reactivity and interactions, especially in biochemical processes.

Resonance in Sulfite

In creating resonance structures for the sulfite ion, we observe that the double bond shared by sulfur and one of the oxygen atoms can be rotated among the oxygen atoms. This shows that the negative charge and the pi electrons are spread over the entire ion, which is why resonance structures play an instrumental role in depicting the stability and the true nature of ions such as the sulfite ion.