Question

Identify each pair that could form a buffer. (a) \(\mathrm{HCl}\) and \(\mathrm{CH}_{3} \mathrm{COOH}\) (b) \(\mathrm{NaH}_{2} \mathrm{PO}_{4}\) and \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\) (c) \(\mathrm{H}_{2} \mathrm{CO}_{3}\) and \(\mathrm{NaHCO}_{3}\)

Short answer

Pairs (b) and (c) can form buffers.

Step-by-step solution

Understand the Concept of a Buffer

A buffer solution can resist changes in pH when small amounts of acid or base are added. Buffers are typically made from a weak acid and its conjugate base, or a weak base and its conjugate acid.

Analyze Pair (a)

The pair given is \(\mathrm{HCl}\) and \(\mathrm{CH}_{3}\mathrm{COOH}\). \(\mathrm{HCl}\) is a strong acid and \(\mathrm{CH}_{3}\mathrm{COOH}\) is a weak acid. A strong acid and a weak acid cannot form a buffer system together because a buffer requires a conjugate acid-base pair.

Analyze Pair (b)

The pair given is \(\mathrm{NaH}_{2}\mathrm{PO}_{4}\) and \(\mathrm{Na}_{2}\mathrm{HPO}_{4}\). \(\mathrm{NaH}_{2}\mathrm{PO}_{4}\) is the conjugate acid and \(\mathrm{Na}_{2}\mathrm{HPO}_{4}\) is the conjugate base. This pair can form a buffer because they constitute a weak acid (\(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\)) and its conjugate base \(\mathrm{HPO}_{4}^{2-}\).

Analyze Pair (c)

The pair given is \(\mathrm{H}_{2}\mathrm{CO}_{3}\) and \(\mathrm{NaHCO}_{3}\). \(\mathrm{H}_{2}\mathrm{CO}_{3}\) is the weak acid and \(\mathrm{NaHCO}_{3}\) provides the conjugate base \(\mathrm{HCO}_{3}^{-}\). This combination forms a buffer system made from a weak acid and its conjugate base.

Key concepts

Understanding Weak Acids

A weak acid is an acid that does not completely dissociate in water. This means that in a solution, it exists in equilibrium with its deprotonated form, which is its conjugate base.
Due to this partial dissociation, weak acids have a higher pH than strong acids. Key characteristics of weak acids include:

• Partial ionization in water
• Establishment of equilibrium between the acid and its conjugate base
• Weaker electrical conductivity compared to strong acids

A classic example of a weak acid is acetic acid (\(\mathrm{CH}_{3}\mathrm{COOH}\)), which partially ionizes in water to produce \(\mathrm{CH}_{3}\mathrm{COO}^{-}\) (the conjugate base) and \(\mathrm{H}^{+}\) ions. The presence of both \(\mathrm{CH}_{3}\mathrm{COOH}\) and \(\mathrm{CH}_{3}\mathrm{COO}^{-}\) in solution allows for the buffering action. This equilibrium is crucial for the function of buffer solutions.

Role of Conjugate Base

A conjugate base is formed when an acid donates a proton (\(\mathrm{H}^{+}\)) to the solution. This base can then accept a proton back, allowing the acid-base equilibrium to occur, which is essential for buffering.
Conjugate bases are important because they help maintain the pH of a solution when small amounts of acid or base are added.In the context of a buffer solution, the presence of both the weak acid and its conjugate base allows the solution to neutralize added acids or bases, thus resisting large changes in pH.

• Example: \(\mathrm{H}_{2}\mathrm{CO}_{3}\) (carbonic acid) ), which turns into its conjugate base, \(\mathrm{HCO}_{3}^{-}\) (bicarbonate).
• In a buffer, if you add more hydrogen ions (\(\mathrm{H}^{+}\)), they will react with the conjugate base. Similarly, if hydroxide ions (\(\mathrm{OH}^{-}\)) are added, they can be neutralized by the weak acid.

Understanding the role of conjugate bases helps in explaining why certain combinations, like \(\mathrm{H}_{2}\mathrm{CO}_{3}\) and \(\mathrm{NaHCO}_{3}\), form effective buffer systems.

How Buffers Show pH Resistance

Buffer solutions are special because they maintain their pH even if small amounts of acid or base are added. This "pH resistance" is due to the equilibrium that exists between the weak acid and its conjugate base in the solution.
When an acid or base is introduced to a buffer, it reacts with either the weak acid or the conjugate base, minimizing changes to the pH.Here's how buffers work:

• If you add an acid, the conjugate base in the buffer will neutralize the added \(\mathrm{H}^{+}\) ions.
• If you add a base, the weak acid in the buffer will donate \(\mathrm{H}^{+}\) ions to neutralize the added \(\mathrm{OH}^{-}\) ions.

This process allows buffers to protect the solution from drastic pH changes. For example, in a \(\mathrm{NaH}_{2}\mathrm{PO}_{4}\) and \(\mathrm{Na}_{2}\mathrm{HPO}_{4}\) buffer system, phosphate ions act to maintain the pH by taking up extra \(\mathrm{H}^{+}\) or \(\mathrm{OH}^{-}\). Understanding this mechanism is key to recognizing why certain pairs, like the phosphate and carbonate buffers, are used in various chemical applications.