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) A localized π bond is a covalent bond formed by the overlapping of two adjacent p orbitals and is restricted to two specific atoms, while a delocalized π bond involves the overlapping of p orbitals on more than two atoms, decreasing overall energy and increasing stability. (b) To determine delocalized π bonding: identify atoms with available p orbitals, determine atom hybridizations, check for resonance structures, and inspect electron distribution. (c) In the nitrite ion (NO₂⁻), the π bond is delocalized as the π electrons are distributed across both oxygen atoms, and there are resonance structures that show the π bond alternating between the two oxygen atoms.

Step-by-step solution

(a) Difference between localized and delocalized π bonds

A localized π bond is a covalent bond formed by the overlapping of two adjacent p orbitals and is restricted to two specific atoms. This type of π bond is found in simple molecules like ethene (C₂H₄). A delocalized π bond, on the other hand, involves the overlapping of p orbitals on more than two atoms. The electrons in such bonds are effectively spread out over several atoms within the molecule, essentially lowering the overall energy and increasing the stability of the system. This is commonly seen in compounds like benzene (C₆H₆), where the π electrons are delocalized over all six carbon atoms.

(b) Determine delocalized π bonding in a molecule or ion

In order to identify if a molecule or ion exhibits delocalized π bonding, we can follow these steps: 1. Identify the atoms that have p orbitals available for π bonding. This typically includes carbon, nitrogen, and oxygen atoms. 2. Determine the hybridization of the atoms identified in step 1. Atoms that are sp² or sp hybridized can participate in delocalized π bonding. 3. Check for resonance structures. If a molecule or ion has multiple resonance structures, it is likely to have delocalized π bonding. 4. Inspect the electron distribution across the molecule or ion. If the π electrons are distributed across multiple atoms within the molecule, then the π bonding is delocalized.

(c) π bond in nitrite ion (NO₂⁻) - localized or delocalized?

We will now determine whether the π bond in NO₂⁻ is localized or delocalized by following the steps from part (b): 1. The nitrite ion, NO₂⁻, contains one nitrogen atom and two oxygen atoms, all of which have p orbitals available for π bonding. 2. The nitrogen atom in NO₂⁻ is sp² hybridized, while the oxygen atoms are sp³ hybridized. 3. NO₂⁻ has resonance structures in which the π bond alternates between the two oxygen atoms. 4. The π electrons in NO₂⁻ are delocalized across both oxygen atoms. Based on this analysis, the π bond in the nitrite ion, NO₂⁻, is delocalized.

Key concepts

Localized π Bond

Localized π bonds are a fundamental notion in molecular chemistry. They represent a simple form of covalent bonding where the π electrons are confined between two specific atoms. This restriction occurs due to the overlap of two adjacent p orbitals, meaning these bond electrons do not extend beyond the two involved atoms.

This kind of bonding is typically observed in smaller or simpler molecules. A classic example would be ethene ( C_2H_4 ), where each pair of pi electrons is tightly bound between two carbon atoms. Due to their localized nature, these bonds do not contribute to the overall aromatic stability often observed with large, delocalized electron systems.

Localized π bonds play a crucial role in the structural and energy profile of the molecules to which they belong, generally resulting in lower stability compared to molecules containing delocalized π bonds.

Molecular Resonance

Molecular resonance is a concept that describes how certain molecules can be represented by two or more equally possible structures. These structures, known as resonance structures, contribute to the true form of the molecule, which is a hybrid of these forms.

Resonance significantly influences the distribution of electrons across a molecule. It typically occurs when there are multiple positions for π bonds to form. A molecule that demonstrates resonance tends to have more delocalized electrons, which contributes to its stability.

• Example: Benzene ( C_6H_6 ) is a well-known molecule that exhibits resonance. It has alternating single and double bonds in a cyclical structure, allowing the bonding electrons to be delocalized around the ring.
• Another illustration is the nitrite ion NO₂⁻, which forms resonance structures by alternating double bonds between the nitrogen and oxygen atoms. This results in a hybridized molecule where the π electrons are shared across the nitrite ion.

By considering resonance structures, chemists can more accurately predict molecular behavior, stability, and reactivity.

Hybridization

Hybridization is a chemical concept used to explain how atomic orbitals fuse to form new hybridized orbitals, which impacts the geometry and bonding properties of molecules.

For atoms that can partake in π bonding, such as carbon, nitrogen, and oxygen, their hybridization state dictates whether they can form localized or delocalized π bonds. Generally, carbon atoms are sp² or sp hybridized when participating in delocalized π bonding. Sp³ hybridized atoms, on the other hand, typically do not partake in π bonding as their orbitals are fully engaged in σ bonds.

• The sp² hybridization involves one s orbital and two p orbitals, forming three hybrid orbitals and leaving a p orbital free for π bonding. This arrangement supports the formation of delocalized π bonds.
• In the case of the sp hybridization, one s orbital combines with one p orbital, resulting in two hybrid orbitals with two p orbitals available for forming π bonds.

Understanding hybridization aids in predicting the 3D shape and bonding characteristics of molecules, which are crucial for grasping their chemical properties and interactions.