Question

Which of the following molecules does not show any resonating structures? (a) \(\mathrm{NH}_{3}\) (b) \(\mathrm{CO}_{3}^{2-}\) (c) \(\mathrm{O}_{3}\) (d) \(\mathrm{SO}_{3}\)

Short answer

\(\mathrm{NH}_{3}\) does not show any resonating structures.

Step-by-step solution

Understand Resonance

Resonance in chemistry is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by a single Lewis formula. A molecule exhibits resonance when it can be represented by two or more valid Lewis structures that have the same arrangement of atoms but different arrangements of electrons.

Examine each molecule for possible resonance structures

Analyse each of the given compounds to determine if it can have different valid Lewis structures with the same arrangement of atoms but different electron configurations.

Evaluate \(\mathrm{NH}_{3}\)

Ammonia (\(\mathrm{NH}_{3}\)) has a single nitrogen atom bonded to three hydrogen atoms and has one lone pair of electrons on the nitrogen. It cannot have resonance as there are no pi bonds or delocalized electrons that can lead to alternative arrangements.

Evaluate \(\mathrm{CO}_{3}^{2-}\)

Carbonate (\(\mathrm{CO}_{3}^{2-}\)) has three oxygen atoms bonded to a central carbon atom and shows resonance. There are multiple ways to draw the structure with the double bond placed between the carbon and any of the three oxygen atoms.

Evaluate \(\mathrm{O}_{3}\)

Ozone (\(\mathrm{O}_{3}\)) also shows resonance, as the double bond can be located between the central oxygen and either of the two outer oxygen atoms, allowing for multiple Lewis structures.

Evaluate \(\mathrm{SO}_{3}\)

Sulfur trioxide (\(\mathrm{SO}_{3}\)) has a sulfur atom surrounded by three oxygen atoms and exhibits resonance as the double bond can be moved among the three sulfur-oxygen bonds.

Determine the molecule without resonance

Based on the evaluation, ammonia (\(\mathrm{NH}_{3}\)) is the molecule that does not show any resonating structures as it lacks delocalized electrons that could lead to alternative Lewis structures.

Key concepts

Lewis Structures

Lewis structures, also known as Lewis dot diagrams, are diagrams that represent the valence electrons of atoms within a molecule. These structures are pivotal for understanding molecular bonding and the arrangement of electrons. To draw a Lewis structure, start by identifying the number of valence electrons each atom contributes to the molecule, then arrange the atoms to satisfy the octet rule, ensuring that each non-hydrogen atom has eight electrons in its valence shell.

Let's take water (H_2O) as an example. Oxygen has six valence electrons, while each hydrogen has one. The oxygen atom is placed in the center, with two hydrogen atoms on the sides connected by single bonds (representing two valence electrons). The remaining four valence electrons of oxygen are indicated as lone pairs on the oxygen atom.

In our exercise, ammonia (NH_3) follows similar principles but ends up with a stable configuration without alternative resonance structures, which clearly differentiates it from the other molecules listed that have multiple valid Lewis structures.

Delocalized Electrons

Delocalized electrons are not associated with a single atom or bond within a molecule but spread out across multiple atoms. This concept is essential in understanding molecules that exhibit resonance. Delocalization occurs in systems containing pi bonds or with conjugated double bonds, where electrons can move through the p-orbitals of neighboring atoms.

For instance, benzene, with its six carbon atoms, exhibits electron delocalization through its conjugated pi system. Similarly, CO_3^{2-} and SO_3 from the given exercise have delocalized electrons that allow for different resonating structures. These pi electrons are not confined to a single bond or location but can be represented in multiple valid ways through resonance. The presence of delocalized electrons can lower the potential energy of the molecule, making the molecule more stable.

Resonating Structures

Resonating structures, often referred to as resonance forms, are multiple valid Lewis structures that a single molecule can adopt, illustrating different possible arrangements of electrons. They are a tool chemists use to represent the actual electronic structure of a molecule more accurately when a single Lewis structure does not suffice. Resonance is depicted using double-headed arrows between the different forms.

The carbonate ion, CO_3^{2-}, is a classic example, where any of its three oxygen atoms can be double-bonded to the central carbon atom, while the actual structure is a hybrid of these resonating structures. Each Lewis structure represents a resonance form, and the true form of the molecule is an average, or a resonance hybrid, of these structures. These forms do not represent distinct configurations that the molecule oscillates between; instead, they are conceptual representations of a single, unchanging structure with delocalized electrons.

In our exercise, sulfur trioxide (SO_3) and ozone (O_3) also exhibit resonance. However, ammonia (NH_3), does not exhibit resonance and can be accurately depicted with a single Lewis structure as it lacks pi bonds required for electron delocalization.