Problem 40 - Chemistry and Chemical Reactivity

Question

Write an equation that describes the equilibrium that exists when the weak acid benzoic acid \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CO}_{2} \mathrm{H}\right)\) dissolves in water. Identify each of the four species in solution as either Bronsted acids or Bronsted bases. Does the equilibrium favor the products or the reactants? (In acting as an acid, the \(\left.-\mathrm{CO}_{2} \mathrm{H} \text { group supplies } \mathrm{H}^{+} \text {to form } \mathrm{H}_{3} \mathrm{O}^{+} .\right)\)

Step-by-Step Solution

Answer

C₆H₅CO₂H is an acid, H₂O is a base, C₆H₅CO₂⁻ is a base, and H₃O⁺ is an acid. The equilibrium favors the reactants.

Step 1: Write the Dissociation Equation

Benzoic acid, when dissolved in water, donates a proton to a water molecule. This process can be described by the equilibrium reaction: \[ \text{C}_6\text{H}_5\text{CO}_2\text{H} + \text{H}_2\text{O} \rightleftharpoons \text{C}_6\text{H}_5\text{CO}_2^- + \text{H}_3\text{O}^+ \] In this equation, benzoic acid is the proton donor, the water molecule is a proton acceptor, and the products of the dissociation are the benzoate ion and the hydronium ion.

Step 2: Identify the Bronsted Acids and Bases

In the dissociation equation:- Benzoic acid (\( \text{C}_6\text{H}_5\text{CO}_2\text{H} \)) acts as a Bronsted acid since it donates a proton.- Water (\( \text{H}_2\text{O} \)) acts as a Bronsted base as it accepts a proton.- The benzoate ion (\( \text{C}_6\text{H}_5\text{CO}_2^- \)) is the conjugate base of benzoic acid.- The hydronium ion (\( \text{H}_3\text{O}^+ \)) is the conjugate acid of water.

Step 3: Determine Equilibrium Favoring

The equilibrium of weak acid dissociation reactions, such as benzoic acid, typically favors the reactants. This is because weak acids only partially dissociate in water.

Key Concepts

Bronsted Acids and Bases

The Bronsted-Lowry theory elegantly describes acids and bases in terms of proton transfer. In this theory, an acid is a substance that can donate a proton (\( \text{H}^+ \)), while a base is one that can accept a proton. Such interactions are commonplace in aqueous solutions, where water often acts as a medium.Consider benzoic acid, \( \text{C}_6\text{H}_5\text{CO}_2\text{H} \), which donates a proton to water when it dissolves. Here, benzoic acid acts as a Bronsted acid, undergoing a transformation where its \( \text{CO}_2\text{H} \) group donates an \( \text{H}^+ \) ion to water. This transforms water into a hydronium ion, \( \text{H}_3\text{O}^+ \), solidifying its role as a Bronsted base, accepting the proton from benzoic acid.

• The substance donating the \( \text{H}^+ \) is the Bronsted acid.
• The substance accepting the \( \text{H}^+ \) becomes the Bronsted base.

Each of these processes creates a conjugate pair, emphasizing the reversible nature of acid-base reactions.

Benzoic Acid Dissociation

Benzoic acid dissociation illustrates the subtle yet profound concept of equilibrium in chemical reactions. When benzoic acid dissolves in water, it partially dissociates into a benzoate ion (\( \text{C}_6\text{H}_5\text{CO}_2^- \)) and a hydronium ion (\( \text{H}_3\text{O}^+ \)). This dissociation can be represented by the equilibrium reaction:\[ \text{C}_6\text{H}_5\text{CO}_2\text{H} + \text{H}_2\text{O} \rightleftharpoons \text{C}_6\text{H}_5\text{CO}_2^- + \text{H}_3\text{O}^+ \]

• This reaction shows the conversion of a weak acid into its conjugate base and the ionization of water into its conjugate acid.
• Benzoic acid is the proton donor, making it a Bronsted acid, while water becomes the proton acceptor, acting as a Bronsted base.
• The newly formed benzoate ion is the conjugate base of benzoic acid, completing the transition of acid to base within the dissociation.

This equilibrium showcases the dynamic nature of chemical reactions, where molecules continuously interact and transform, yet ultimately maintain a balance between reactants and products.

Weak Acids

Weak acids like benzoic acid do not completely dissociate in water, producing a mixture of acid and its dissociated ions at equilibrium. Unlike strong acids, which fully dissociate, weak acids only partially release their protons. This incomplete dissociation is crucial for understanding the behavior of such acids in solution.

• Weak acids have higher \( \text{pK}_a \) values, indicating their lower degree of ionization compared to strong acids.
• In the case of benzoic acid, the equilibrium heavily favors the reactant side, with most of the acid molecules remaining undissociated in solution.
• This partial dissociation results in a relatively low concentration of hydronium ions, which means weak acids are less effective at lowering the pH of a solution.

This understanding underscores the importance of acid strength and equilibrium concepts in predicting the results of chemical reactions. Weak acids are pivotal in many natural processes and industrial applications, providing less aggressive, controlled reactions.