Question

I am struggling with what appears to be an extremely easy pH problem that uses the Henderson-Hasselbalch equation. The problem and answer provided by the book is given below. I understand where the book is getting the answer but it appears that they are overlooking something. From what I understand, the Henderson-Hasselbalch equation allows you to calculate the pH of a solution based on the concentrations of acid and conjugate base at equilibrium. In the problem, you start off with a solution of just the conjugate base. Then you add pure acid to the solution to get a desired pH. However, the book seems to assume that the acid you add completely contributes to the concentration of undissociated acid at equilibrium [HA]. How can this be the case?

Answer 1

The Henderson-Hasselbalch equation approximates the pH of buffer solutions based on the relation of weak acid and conjugate base concentrations. It presumes complete reaction of added strong acid with the conjugate base, affecting the equilibrium concentrations of acid and base.

Step-by-step explanation:

The Henderson-Hasselbalch equation is an approximation used to calculate the pH of buffer solutions, which involves the concentration of a weak acid (HA) and its conjugate base (A-).

The equation is pH = pKa + log([A-]/[HA]), based on the assumption that the acid-base reaction is at equilibrium. In the case of adding pure acid to a solution, the premise is that the acid (assumed to be strong) will react completely with the conjugate base present in the buffer, transforming it into the weak acid HA, therefore contributing to the [HA] at equilibrium without significantly altering the pH due to the buffering action.

However, this assumption may not always hold true, especially if the buffer capacity is exceeded, or the acid added is weak and doesn't dissociate completely. For accurate results, one must consider the extent of the reaction between the acid and the conjugate base and the consequent shift in equilibrium.