Question

Calculate formal charges on the \(\mathrm{C}\) and \(\mathrm{O}\) atoms in two resonance structures for acetic acid, the main component of vinegar. Decide which structure makes a larger contribution to the resonance hybrid.

Short answer

The first structure contributes more to the resonance hybrid.

Step-by-step solution

Draw the Lewis Structures

First, we need to draw the two main resonance structures for acetic acid (CH₃COOH). The first structure shows the carbon double-bonded to oxygen (C=O) with a single bond to hydroxyl group (OH). In the second structure, consider the possibility of moving the electrons to show another resonance form (such as an enolate form).

Calculate Formal Charge for Carbon

To find the formal charge, use the formula: Formal charge = (Valence electrons) - (Non-bonding electrons) - 0.5*(Bonding electrons). For the carbon atom double-bonded to oxygen in the first structure: - Carbon has 4 valence electrons, 0 lone pairs, and 4 bonding pairs. - Formal charge = 4 - 0 - (4/2) = 0.

Calculate Formal Charge for Oxygen in First Structure

For the double-bonded oxygen (C=O): - Oxygen has 6 valence electrons, 4 non-bonding electrons, and 2 bonding pairs. - Formal charge = 6 - 4 - (2/2) = 0. For the single-bonded oxygen (OH group): - Oxygen has 6 valence electrons, 6 non-bonding electrons, and 1 bonding pair. - Formal charge = 6 - 6 - (2/2) = -1.

Calculate Formal Charges in Second Resonance Structure

In the second structure, assume that the carbon has a single bond to the oxygen (C-O) and that oxygen has taken a lone pair from carbon, resembling an enolate. For the carbon now singly bonded to oxygen: - Formal charge = 4 - 0 - (6/2) = +1. For the double-bonded oxygen, the formal charge remains 0, and for the single-bonded oxygen: - Formal charge = 6 - 6 - 1 = -1.

Compare the Resonance Structures

The first resonance structure, having formal charges of 0 for the carbon and 0 and -1 for oxygens, is more stable due to minimized charges and the negative charge on the more electronegative oxygen. The second structure, bearing a positive charge on carbon, is less favorable energetically.

Key concepts

Resonance Structures

In chemistry, resonance structures are essential for understanding the stability and behavior of certain molecules. They offer various ways to represent the electron distribution in a molecule without changing the actual arrangement of atoms. This concept is crucial because some molecules can't be described by a single Lewis structure satisfactorily.

Resonance structures illustrate how electrons can be shifted within a molecule. By considering different resonance structures, chemists can better predict the most stable electron configuration. These structures don't exist independently in reality but are used to visualize potential states of electron distribution. The true representation of the molecule is called a resonance hybrid, which is a weighted average of all possible resonance structures. In acetic acid, for instance, resonance helps us understand the electron shifts between carbon, oxygen, and hydrogen atoms. This process accounts for the molecule’s stability and reactivity, especially when analyzing the effects of various resonance contributors on stability.

Lewis Structures

Lewis structures are diagrams that show the bonding between atoms and the lone pairs of electrons in a molecule. They are a way to represent molecules visually by depicting electron distribution around the atoms. In Lewis structures, we make use of dots to show lone electrons and lines to represent shared electron pairs that form a chemical bond.

The creation of Lewis structures involves a set of steps that begin with counting the total number of valence electrons in the molecule. Once you've determined this number, you arrange the atoms in a way that satisfies the octet rule as much as possible. The goal is to depict each atom reaching a stable electronic configuration by sharing or transferring electrons. Acetic acid, being an organic molecule, contains both single and double bonds that are important to feature accurately in its Lewis structure with respect to its resonance forms. By examining the Lewis structures of acetic acid's resonance forms, we can calculate formal charges and make predictions about molecular stability.

Acetic Acid

Acetic acid, commonly known as the primary component of vinegar, is represented by the chemical formula \( ext{CH}_3 ext{COOH}\). It is an important molecule in organic chemistry and everyday life. Structurally, acetic acid consists of a methyl group (\( ext{CH}_3\)) linked to a carboxylic acid group (\( ext{COOH}\)), making it a simple carboxylic acid.

Understanding the structure of acetic acid involves looking at the connections and bonds between its atomic components. The carbon atoms form the backbone, with the presence of a carbonyl group (a carbon double bonded to an oxygen, \( ext{C}=\text{O}\)) adjacent to the hydroxyl group (\( ext{OH}\)). These functional groups are significant because they contribute greatly to the chemical properties of acetic acid, including acidity and its ability to participate in chemical reactions through the exchange of protons.

By using resonance structures and Lewis structures, we can better appreciate how acetic acid exists in resonance forms, influencing its stability and reactivity. It's this understanding that allows chemists to predict and explain the diverse interactions and behaviors of acetic acid in both natural and industrial contexts.