Question

(a) What is the difference between a localized \(\pi\) bond and a delocalized one? (b) How can you determine whether a molecule or ion will exhibit delocalized \(\pi\) bonding? (c) Is the \(\pi\) bond in \(\mathrm{NO}_{2}^{-}\) localized or delocalized?

Short answer

A localized π bond is confined between two adjacent atoms, while a delocalized π bond involves electrons spread over more than two nuclei in conjugated systems or aromatic compounds. To identify delocalized π bonding, look for conjugated double bonds, aromaticity, and resonance structures. The π bond in NO₂⁻ is delocalized due to the presence of resonance structures, where the charge and bonding electrons are distributed across all three atoms.

Step-by-step solution

Part A: Localized and Delocalized π Bonds

A localized π bond is a covalent bond formed by the sideways overlap of p orbitals. In a localized π bond, the bonding electrons are confined between two adjacent atoms. On the other hand, a delocalized π bond is a bond where electrons are spread over more than two nuclei, creating an extended π system. This extended π system is typically observed in conjugated systems or aromatic compounds, where the electrons are spread over several atoms and not limited to two specific atoms.

Part B: Identifying Delocalized π Bonding

To determine if a molecule or ion has delocalized π bonding, we can look for the following features: 1. The presence of conjugated double bonds: A system with alternating single and double bonds allows for the π electrons to be delocalized across multiple bond sites. 2. Aromaticity: A compound that follows the Hückel's rule (4n+2 π electrons, where n is an integer) may have delocalized π bonding, as seen in compounds like benzene. 3. Resonance structures: If a molecule or ion has multiple valid Lewis structures with different electron arrangements, this can be an indication of delocalized π bonding, as the electrons are distributed across multiple sites.

Part C: π Bond in NO₂⁻

To determine if the π bond in NO₂⁻ is localized or delocalized, we need to examine the resonance structures. 1. Draw the Lewis structure of NO₂⁻: The central atom is nitrogen (N) with two oxygen atoms (O) bonded to it, and one extra electron from the negative charge. 2. Determine resonance structures: The nitrogen atom forms a double bond with one of the oxygen atoms and a single bond with the other, with a lone electron pair on the single-bonded oxygen atom. This structure can resonate with the other oxygen atom to form another possible structure with equal stability. Since both structures are equivalent, the real structure is a resonance hybrid, where the charge and bonding electrons are delocalized across the three atoms. 3. Delocalization decision: Due to the resonance phenomenon between the two possible structures, the π bond in NO₂⁻ is delocalized.