Question

Which of the following solutions is a buffer? (a) 0.10\(M\) \(\mathrm{CH}_{3} \mathrm{COOH}\) and \(0.10 \mathrm{MCH}_{3} \mathrm{CONa},(\mathbf{b}) 0.10 \mathrm{MCH}_{3} \mathrm{COOH}\) (c) 0.10 \(\mathrm{M} \mathrm{HCl}\) and \(0.10 \mathrm{M} \mathrm{NaCl},(\mathbf{d})\) both a and \(\mathrm{c},(\mathbf{e})\) all of a, \(\mathrm{b},\) and \(\mathrm{c} .\)

Short answer

The correct answer is (a) 0.10 M CH3COOH and 0.10 M CH3COONa, as it consists of a weak acid (acetic acid) and its conjugate base (acetate).

Step-by-step solution

Identify the components of each solution

In the given exercise, we are given the following solutions: (a) 0.10 M CH3COOH and 0.10 M CH3COONa (b) 0.10 M CH3COOH (c) 0.10 M HCl and 0.10 M NaCl (d) Both a and c (e) All of a, b, and c We now need to identify the components of each solution and determine if they can form a buffer.

Analyze the solutions

(a) 0.10 M CH3COOH (acetic acid) and 0.10 M CH3COONa (sodium acetate): In this solution, CH3COOH is a weak acid and CH3COO- (from the dissociation of CH3COONa) is its conjugate base. Thus, this solution can act as a buffer. (b) 0.10 M CH3COOH (acetic acid): In this solution, there is only a weak acid present, and no conjugate base is available. Therefore, this solution cannot act as a buffer. (c) 0.10 M HCl (hydrochloric acid) and 0.10 M NaCl (sodium chloride): In this solution, HCl is a strong acid and completely dissociates into H+ and Cl-. Since HCl is not a weak acid, this solution cannot be a buffer.

Choose the correct answer

Based on our analysis in step 2, we can conclude that: (a) is a buffer (b) is not a buffer (c) is not a buffer Thus, the correct answer is (a) 0.10 M CH3COOH and 0.10 M CH3COONa.

Key concepts

Acid-Base Equilibrium

Acid-base equilibrium is a fundamental concept in chemistry that deals with the balance between acids and bases in a solution. This balance is crucial for buffer solutions, which maintain a stable pH despite the addition of small amounts of acid or base.
To understand acid-base equilibrium, it is essential to know that acids donate protons (\( H^+ \)) while bases accept them. For a buffer, this equilibrium is achieved by having a weak acid and its conjugate base present in solution. The weak acid can donate \( H^+ \), while the conjugate base can accept \( H^+ \), thus resisting changes in pH.
A classic example of acid-base equilibrium is seen in the acetic acid (\( ext{CH}_3 ext{COOH} \)) and sodium acetate (\( ext{CH}_3 ext{COONa} \)) buffer system. Here, acetic acid partially dissociates to produce its conjugate base acetate (\( ext{CH}_3 ext{COO}^- \)), which maintains equilibrium by reacting with added acids or bases.

• If a strong acid is added, the acetate ions react with the \( H^+ \) ions, neutralizing them.
• If a strong base is added, the acetic acid donates \( H^+ \) to neutralize the hydroxide ions (\( ext{OH}^- \)).

This balance ensures that the pH of the solution remains relatively constant, which is vital for many chemical processes and biological functions.

Conjugate Acid-Base Pairs

Conjugate acid-base pairs are pairs of molecules or ions that are connected through the gain or loss of a proton (\( H^+ \)). In any acid-base reaction, an acid is transformed into its conjugate base, while a base is transformed into its conjugate acid.
For example, in the acetic acid-sodium acetate buffer system, acetic acid (\( ext{CH}_3 ext{COOH} \)) acts as the acid. When it donates a proton, it forms its conjugate base, acetate (\( ext{CH}_3 ext{COO}^- \)). Conversely, acetate can act as a base by accepting a proton to reform acetic acid.
Understanding conjugate acid-base pairs is crucial in recognizing how buffers stabilize pH. When a buffer solution is disturbed by the addition of a strong acid or base:

• The conjugate base absorbs excess \( H^+ \) ions from strong acids, preventing significant pH changes.
• The conjugate acid provides \( H^+ \) ions if a strong base is added, again stabilizing the pH.

This dynamic operation of conjugate pairs ensures that buffers effectively moderate pH changes, making them invaluable in both laboratory and biological contexts.

Weak Acids and Strong Acids

Weak and strong acids differ in their ability to dissociate in water. This distinction is key in understanding their roles in buffers. Weak acids, such as acetic acid, do not fully dissociate in water, which allows them to play a crucial role in buffer systems.
Weak Acids:

• Partially dissociate in solution.
• Exist in equilibrium with their conjugate base.
• Suitable for buffers as their dissociation is reversible, allowing equilibrium to be maintained.

Strong acids, such as hydrochloric acid (\( ext{HCl} \)), behave differently. They fully dissociate in water, releasing all their protons.

• This complete dissociation means they can't form effective buffers because there is no reverse reaction to restore equilibrium.

The inefficacy of strong acids in buffer solutions is evident in the fact that they do not provide the equilibrium stability that weak acids and their conjugate bases do. Thus, understanding the role of weak and strong acids is fundamental in selecting components for creating an effective buffer system.