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 bonds are formed between two adjacent atoms and are restricted to the region between them, whereas delocalized π bonds involve overlapping of more than two atomic orbitals across three or more atoms. (b) To determine if a molecule or ion exhibits delocalized π bonding, look for resonance structures with π bonds shared across three or more atoms. (c) The π bond in the nitrite ion, \(\mathrm{NO}_{2}^{-}\), is delocalized.

Step-by-step solution

Explaining Localized and Delocalized π bonds

Localized π bonds are formed between two adjacent atoms and are restricted to the region between these two atoms. In contrast, delocalized π bonds involve the overlapping of more than two atomic orbitals across three or more atoms. This leads to an electron cloud that extends over multiple atoms, distributing the bonding electrons across a larger area.

Determination of Delocalized π bonds in a Molecule or Ion

To determine whether a molecule or ion exhibits delocalized π bonding, examine the structure of the compound and look for resonance structures. If there is more than one possible structure with π bonds, check whether these structures share electrons across three or more atoms. If the electrons are distributed over multiple atoms, this suggests delocalized π bonding.

Analyzing the π bond in the Nitrite Ion, \(\mathrm{NO}_{2}^{-}\)

The nitrite ion (\(\mathrm{NO}_{2}^{-}\)) has two resonance structures, as shown below: (1) O-N=O^(-) ↔ (2) O^(-)-N=O In these resonance structures, the π bond is distributed between the two oxygen atoms and the nitrogen atom. The electron cloud that forms the π bond is spread over three atoms, which indicates delocalized π bonding.

Answer

(a) Localized π bonds are formed between two adjacent atoms and are restricted to the region between them, whereas delocalized π bonds involve overlapping of more than two atomic orbitals across three or more atoms. (b) To determine if a molecule or ion exhibits delocalized π bonding, look for resonance structures with π bonds shared across three or more atoms. (c) The π bond in the nitrite ion, \(\mathrm{NO}_{2}^{-}\), is delocalized.

Key concepts

Localized π Bonds

Localized π bonds are a type of covalent bond where the bonding electrons are strictly confined to the space between two bonded atoms. This means that these electrons are not shared with or over any other atom within the molecule. In cases with localized π bonds, there is a clear and direct electronic interaction between a pair of atoms.

This situation occurs in simple pi-bonded systems, like the double bonds seen in alkenes. Here, the π bond is formed from the sideways overlap of p orbitals only between the two carbon atoms. As a result, the electron density resulting from this overlap is tightly held between these two specific atoms and does not extend to other atoms. This is why such bonds are termed as 'localized.'

Therefore, in a localized π bond scenario, the bond's properties (such as strength and energy) rely heavily on the relationship and distance between two specific atoms.

Resonance Structures

Resonance structures are multiple possible structures for a molecule that depict different arrangements of electrons around the atoms. These structures are crucial for understanding molecules where electrons can be distributed over more than just two atoms.

While drawing resonance structures, we follow the rule that the position of atoms does not change; rather, only the positions of the electrons alter. This means that while the connectivity of atoms remains constant, you can often represent the same molecule by several contributing structures showing the same number of electron pairs.

To identify if a molecule has resonance structures, look for:

• Multiple potential positions for double or π bonds/li>
• Presence of lone pairs that can form additional bonds/li>
• Overall charge that can be relocated across different atoms

For compounds displaying resonance, like some oxyanions and carboxylic acids, these structures help illustrate the actual electron distribution as a hybrid, showing that electrons are delocalized rather than being fixed in one particular bond.

Nitrite Ion (NO2-)

The nitrite ion, \( ext{NO}_2^−\), is an excellent example of delocalized π bonding due to its resonance structures. It features a nitrogen atom centrally bonded to two oxygen atoms, with a negative charge spread across the ion.

In its two resonance forms, the double-bond alternates between each of the two nitrogen-oxygen pairs, presenting as:

• O-N=O^-
• O^--N=O

These forms demonstrate that while the actual double bond is not fixed on one oxygen, the electrons forming this bond are spread evenly across all three atoms. This electron distribution defines the delocalization in the nitrite ion.

The resonance concept in nitrite ion helps stabilize the molecule, as the negative charge and the π bond shift between the two oxygen atoms. This enhances stability compared to a scenario where one single structure carries a fixed charge or bond arrangement. Thus, in NO2-, the π bond is clearly delocalized, leading to a more stably distributed electronic cloud across the molecule.