A friend asks the following: "Consider a buffered solution made up of the weak acid HA and its salt NaA. If a strong base like NaOH is added, the HA reacts with the OH - to form A Thus the amount of acid (HA) is decreased, and the amount of base \(\left(\mathrm{A}^{-}\right)\) is increased. Analogously, adding HCl to the buffered solution forms more of the acid (HA) by reacting with the base \(\left(\mathrm{A}^{-}\right)\). Thus how can we claim that a buffered solution resists changes in the pH of the solution?" How would you explain buffering to this friend?

A buffered solution is a mixture of a weak acid (HA) and its conjugate base, which is usually a salt (NaA). In this context, the weak acid (HA) donates a proton (H+) to the solution, while the conjugate base (A-) accepts a proton. The ratio of the acid to its conjugate base in the solution determines the pH of the buffered solution. The combination of weak acid and its conjugate base has the ability to maintain a relatively constant pH despite the addition of a small amount of strong acid or base.

When a strong base (like NaOH) is added to the buffered solution, it reacts with the weak acid (HA) in the solution, neutralizing it. The reaction is given by: \[HA + OH^- \rightarrow A^- + H_2O\] As a result, the concentration of the weak acid (HA) decreases, and the concentration of its conjugate base (A-) increases.

Conversely, when a strong acid (like HCl) is added to the buffered solution, it reacts with the conjugate base (A-) present in the solution. The reaction is as follows: \[A^- + H^+ \rightarrow HA\] In this case, the concentration of the conjugate base (A-) decreases, and the concentration of the weak acid (HA) increases.

The buffering action is due to the equilibrium that exists between the weak acid (HA) and its conjugate base (A-) in the solution. When a strong base or acid is added, the equilibrium shifts in a direction to counteract the effect of the added base or acid. In the case of a strong base being added, the equilibrium shifts to the right, removing the excess hydroxide ions (OH-) and maintaining the pH. When a strong acid is added, the equilibrium shifts to the left, removing the excess hydrogen ions (H+) and again preserving the pH.

The pH stability of a buffered solution can be explained using the Henderson-Hasselbalch equation: \[pH = pK_a + \log \frac{[A^-]}{[HA]}\] Here, \(pK_a\) is the negative logarithm of the acid dissociation constant of the weak acid (HA). When the ratio of the conjugate base (A-) to the weak acid (HA) remains relatively constant, the pH of the buffered solution also remains relatively constant. The addition of small amounts of strong acid or base will not significantly change the ratio of the two components, hence the pH is maintained. In conclusion, a buffered solution resists changes in pH by maintaining a constant ratio of weak acid to its conjugate base, even when a strong acid or base is added to the solution. It counteracts the added acid or base by shifting the acid-base equilibrium, thus maintaining the buffered solution's pH.