Problem 61 - Chemistry and Chemical Reactivity

Question

Draw the two resonance structures that describe the bonding in the acetate ion. What is the hybridization of the carbon atom of the \(-\mathrm{CO}_{2}^{-}\) group? Select one of the two resonance structures and identify the orbitals that overlap to form the bonds between carbon and the three elements attached to it.

Step-by-Step Solution

Answer

The acetate ion has two equal resonance structures. Carbon is \(sp^2\) hybridized.

Step 1: Draw the Lewis Structure

Start by drawing the basic Lewis structure for the acetate ion \( ext{CH}_3 ext{CO}_2^−\). The central carbon atom is bonded to three oxygen atoms and one of these oxygen atoms carries a negative charge. The other oxygen atom is double-bonded to the carbon.

Step 2: Identify Resonance Structures

The acetate ion has two resonance structures. In one structure, the double-bond is formed between carbon and one oxygen, while in the second resonance structure, the double bond between carbon and oxygen shifts to the other oxygen atom. Both resonance forms contribute equally to the resonance hybrid.

Step 3: Determine Hybridization of Carbon

In the \(- ext{CO}_2^−\) group of the acetate ion, the central carbon is bonded to three regions (one oxygen with a double bond, the other oxygen with a single bond, and the methyl group). This implies \(sp^2\) hybridization for the carbon as there are three electron domains.

Step 4: Choose a Resonance Structure

Select one of the resonance structures, say the one with a double bond to the left oxygen. Here, identify the bonding: the \( ext{C}- ext{O}\) double bond involves \(sp^2\) hybrid orbitals from carbon overlapping with \(p\) orbitals from oxygen. The single \( ext{C}- ext{O}\) bond involves an \(sp^2\) hybrid orbital from carbon overlapping with another \(sp^3\) hybrid orbital from oxygen.

Step 5: Bonding Overview for Carbon

In the selected resonance structure, all bonds are formed through orbital overlap: the carbon forms one \( ext{C}- ext{O}\) double bond, comprising a sigma bond (\(sp^2-sp^2\) overlap) and a pi bond (\(p-p\) overlap), and a single \( ext{C}- ext{O}\) bond via \(sp^2-sp^3\) overlap. The \( ext{C}- ext{C}\) bond is formed by \(sp^2-sp^3\) overlap as well.

Key Concepts

Resonance Structures

Resonance structures in chemistry are quite fascinating! They help us understand how electrons can be arranged in a molecule. Let's look into the acetate ion as an example. The acetate ion, \\( \text{CH}_3\text{CO}_2^- \), has two important resonance structures.
In these structures, the configuration of atoms stays the same, but the placement of electrons can differ.
Both resonance structures of the acetate ion involve a carbon atom bonded to two oxygen atoms.

• In one structure, a double bond is formed between the carbon atom and one of the oxygen atoms.
• In the other structure, the double bond is found between the carbon atom and the second oxygen atom.

These structures reflect how electrons can shift yet keep the molecule's stability. Both equally contribute to the resonance hybrid, presenting an intermediate form.
Understanding this can help us learn how flexible and resilient molecules really are when it comes to bonding.

Hybridization

Hybridization is a concept that explains the mixing of atomic orbitals to form new hybrid orbitals. This is so crucial because it helps us predict and explain the geometry and bonding properties of molecules.
For the acetate ion, focus on the central carbon atom within the \\( -\text{CO}_2^- \) group. This specific carbon atom is fascinating because it exhibits \( sp^2 \) hybridization.
So, what does \( sp^2 \) hybridization mean?

• \( sp^2 \) hybridization implies that one \( s \) and two \( p \) orbitals mix.
• This results in three \( sp^2 \) hybridized orbitals.
• These hybridized orbitals create a trigonal planar shape with 120-degree bond angles.

In the acetate ion, this configuration suits its bond formation perfectly. The carbon atom accommodates three regions of electron density: two bonds with oxygen atoms and one with the methyl group. Such hybridization is vital for the molecule's stability and function as it determines its shape.

Carbon-Oxygen Bonding

The bonding between carbon and oxygen in the acetate ion is pretty unique due to the molecule's structure. Let's delve deeper into how these bonds form in one specific resonance structure of the ion.
The acetate ion has two types of carbon-oxygen bonds: a double bond and a single bond.
Here's how they differ:

• Double bond (C=O): This consists of one sigma (\( \sigma \)) and one pi (\( \pi \)) bond. The sigma bond occurs from the \( sp^2 \) hybrid orbital of carbon overlapping with the \( sp^2 \) orbital of oxygen. The pi bond forms from the unhybridized \( p \) orbitals overlapping from both atoms.
• Single bond (C-O): Involves a sigma bond resulting from \( sp^2 \) hybridization on carbon overlapping with an \( sp^3 \) hybrid orbital on the corresponding oxygen atom.

Understanding these bonds gives insights into the molecule's reactivity and properties. While the acetate ion resonates between different structures, the nature of these bonds provides foundational stability to the molecule.

Lewis Structure

Lewis structures are like blueprints for molecules. They show how atoms are connected and where electrons are located. For the acetate ion, \( \text{CH}_3\text{CO}_2^- \), drawing a Lewis structure reveals valuable information about its electron configuration and bonding.
To draw a Lewis structure, specific steps should be followed:

• Determine the total number of valence electrons. For the acetate ion, consider both carbon, hydrogen, and oxygen contributions, along with the extra electron from its negative charge.
• Arrange the atoms with carbon bonded to two oxygen atoms and a methyl group.
• Complete the octets around oxygen atoms and distribute any leftover electrons. Ensure all atoms (except hydrogen) achieve an octet configuration.

Lewis structures help not only visualize molecular shapes but also predict reactivity and interactions. For instance, the presence of a lone pair on one of the oxygen atoms often indicates potential sites for chemical reactions in the acetate ion.