Question

Except for the \(\mathrm{Na}^{+}\) spectator ion, aqueous solutions of \(\mathrm{CH}_{3} \mathrm{COOH}\) and \(\mathrm{CH}_{3} \mathrm{COONa}\) contain the same species. (a) What are the species (other than \(\left.\mathrm{H}_{2} \mathrm{O}\right) ?\) (b) Why is \(0.1 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}\) acidic and \(0.1 M \mathrm{CH}_{3} \mathrm{COONa}\) basic?

Short answer

(a) \(\mathrm{CH}_{3} \mathrm{COOH}\), \(\mathrm{CH}_{3} \mathrm{COO}^-\), \(\mathrm{H}^+\), and \(\mathrm{H}_2\mathrm{O}\). (b) \(\mathrm{CH}_{3} \mathrm{COOH}\) is acidic due to \(\mathrm{H}^+\) ions; \(\mathrm{CH}_{3} \mathrm{COONa}\) is basic due to \(\mathrm{CH}_{3} \mathrm{COO}^-\) forming \(\mathrm{OH}^-\) ions.

Step-by-step solution

- Identify species in solutions

First, identify the species present in the aqueous solutions of \(\mathrm{CH}_{3} \mathrm{COOH}\) and \(\mathrm{CH}_{3} \mathrm{COONa}\). \(\mathrm{CH}_{3} \mathrm{COOH}\), being a weak acid, partially dissociates in water forming \(\mathrm{CH}_{3} \mathrm{COO}^-\) and \(\mathrm{H}^+\). \(\mathrm{CH}_{3} \mathrm{COONa}\) dissociates completely in water into \(\mathrm{CH}_{3} \mathrm{COO}^-\) and \(\mathrm{Na}^+\). So, we have: Species in \(\mathrm{CH}_{3} \mathrm{COOH}\): \(\mathrm{CH}_{3} \mathrm{COOH}\), \(\mathrm{CH}_{3} \mathrm{COO}^-\), \(\mathrm{H}^+\), \(\mathrm{H}_2\mathrm{O}\). Species in \(\mathrm{CH}_{3} \mathrm{COONa}\): \(\mathrm{CH}_{3} \mathrm{COO}^-\), \(\mathrm{Na}^+\), \(\mathrm{H}_2\mathrm{O}\).

- Explain acidity of \(\mathrm{CH}_{3} \mathrm{COOH}\)

Explain why \(0.1 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}\) is acidic. Acetic acid \(\mathrm{CH}_{3} \mathrm{COOH}\) is a weak acid that dissociates slightly to give \(\mathrm{H}^+\) ions in solution: \[ \mathrm{CH}_{3} \mathrm{COOH} \rightleftharpoons \mathrm{CH}_{3} \mathrm{COO}^- + \mathrm{H}^+ \] The presence of \(\mathrm{H}^+\) ions makes the solution acidic.

- Explain basicity of \(\mathrm{CH}_{3} \mathrm{COONa}\)

Explain why \(0.1 M \mathrm{CH}_{3} \mathrm{COONa}\) is basic. Sodium acetate \(\mathrm{CH}_{3} \mathrm{COONa}\) dissociates completely to form \(\mathrm{CH}_{3} \mathrm{COO}^-\) and \(\mathrm{Na}^+\). The acetate ion \(\mathrm{CH}_{3} \mathrm{COO}^-\) is a conjugate base of acetic acid and can accept \(\mathrm{H}^+\) ions from water, producing \(\mathrm{OH}^-\) ions in the reaction: \[ \mathrm{CH}_{3} \mathrm{COO}^- + \mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{CH}_{3} \mathrm{COOH} + \mathrm{OH}^- \] The production of \(\mathrm{OH}^-\) ions makes the solution basic.

Key concepts

Weak Acids

A weak acid is an acid that only partially dissociates in water. This means only a small fraction of the acid molecules break apart into ions. Acetic acid (\text{CH}_3\text{COOH}) is a classic example of a weak acid. In an aqueous solution, only some acetic acid molecules release protons (\text{H}^+) and form acetate ions (\text{CH}_3\text{COO}^-).

Here's the dissociation reaction for acetic acid:
\( \text{CH}_3\text{COOH} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}^+ \)

This equilibrium signifies that the reaction goes both ways, meaning that some acetate ions and hydrogen ions recombine to form acetic acid again. Because of this partial dissociation, the concentration of \text{H}^+ ions is relatively low, giving the solution a moderate acidic character. Weak acids are present in various everyday substances like vinegar, fruit juices, and many sour-tasting foods.

Dissociation

Dissociation in chemistry refers to the process where molecules split into smaller particles such as ions. In aqueous solutions, this typically happens when a substance dissolves in water. Dissociation can be complete or partial:

• Complete Dissociation: The substance breaks into its ions entirely. For example, sodium acetate (\text{CH}_3\text{COONa}) dissociates completely to give sodium ions (\text{Na}^+) and acetate ions (\text{CH}_3\text{COO}^-).
• Partial Dissociation: Only a fraction of the substance breaks into ions. Acetic acid (\text{CH}_3\text{COOH}), for instance, partially dissociates in water to form \text{H}^+ and \text{CH}_3\text{COO}^- ions.

The equilibrium for acetic acid can be illustrated as:
\( \text{CH}_3\text{COOH} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}^+ \)
The state of equilibrium means that the forward and reverse reactions balance each other out. It’s this delicate balance that determines the degree of acidity or basicity in the solution. Understanding dissociation helps explain why some solutions have more \text{H}^+ or \text{OH}^- ions, leading to their acidic or basic nature.

Conjugate Base

When an acid donates a proton (\text{H}^+), the species that remain is called its conjugate base. The conjugate base can accept \text{H}^+ ions from water or other donors. This concept is crucial for understanding the behavior of acids and bases in water.

Consider acetic acid (\text{CH}_3\text{COOH}). When it loses a \text{H}^+ ion, it forms its conjugate base, the acetate ion (\text{CH}_3\text{COO}^-):
\( \text{CH}_3\text{COOH} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}^+ \)

The acetate ion (\text{CH}_3\text{COO}^-), being the conjugate base, is capable of accepting a \text{H}^+ ion from water. This results in the formation of hydroxide ions (\text{OH}^-) and regenerates acetic acid:
\( \text{CH}_3\text{COO}^- + \text{H}_2\text{O} \rightleftharpoons \text{CH}_3\text{COOH} + \text{OH}^- \)

This exchange increases the concentration of \text{OH}^- ions in the solution, making it basic. Therefore, sodium acetate (\text{CH}_3\text{COONa}), which provides \text{CH}_3\text{COO}^- ions upon complete dissociation, creates a basic environment. Understanding conjugate acid-base pairs helps explain the pH changes in different solutions and is fundamental in predicting the outcome of chemical reactions.