Question

When a strong acid is added to a buffer, the \(\mathrm{pH}\) changes a little bit. (a) Does the pH increase or decrease? (b) Why does the \(\mathrm{pH}\) change at all? Why doesn't the buffer hold the \(\mathrm{pH}\) constant?

Short answer

When a strong acid is added to a buffer, the pH decreases due to the increased concentration of H+ ions. The buffer can only maintain its pH within a certain buffer capacity, determined by the concentrations of the weak acid and its conjugate base. The pH doesn't remain constant because the buffer's capacity to neutralize the added acid is limited, so as more strong acid is added, the pH of the buffer solution will eventually change.

Step-by-step solution

(a) pH change direction when a strong acid is added

When a strong acid is added to a buffer solution, it will donate protons (H+) to the solution, increasing the concentration of H+ ions. Since pH is defined as: \[\mathrm{pH} = -\log[\ce{H+}]\] As the concentration of H+ increases, the logarithm of [H+] will result in a smaller negative number, which means the pH will decrease.

(b) Explanation for the pH change

A buffer solution is designed to minimize changes in pH upon the addition of small amounts of acid or base. However, the buffer solution can only maintain its pH within a certain buffer capacity, determined by the concentrations of the weak acid and its conjugate base present in the solution. When a strong acid is added, the buffer system works to neutralize the extra protons (H+ ions) donated by the strong acid: \[\ce{HA (aq) + H+ (aq) <=> A- (aq) + H2O (l)}\] Where HA is the weak acid and A- is its conjugate base. The strong acid consumed in this reaction supplies the extra H+ that is neutralized, leading to the formation of more HA. However, the weak acid (HA) must be present in sufficient amounts to neutralize the added strong acid; if the strong acid's concentration overwhelms the buffer capacity, the pH of the solution will change. The pH does not remain constant because the buffer's capacity to neutralize the added acid is limited by the concentration of weak acid and its conjugate base. As the buffer keeps neutralizing the added acid, its ability to counteract further changes in pH diminishes. Therefore, as more strong acid is added, the pH of the buffer solution will eventually change.

Key concepts

Buffer Capacity

Buffer capacity refers to the ability of a buffer solution to resist changes in pH when acids or bases are added. It specifically indicates the quantity of acid or base that the buffer can neutralize before the pH begins to change significantly. It is dependent on two main factors: the concentration of the acid and its conjugate base in the buffer solution, as well as the proximity of the solution's pH to the pKa (acid dissociation constant) of the weak acid component of the buffer.

A buffer with high concentrations of the weak acid and its conjugate base can neutralize more added strong acid or base with minimal pH change, indicating a higher buffer capacity. On the other hand, when the pH is near the weak acid's pKa value, the buffer's capacity is at its maximum, since the acid and base forms are present in similar amounts and can effectively neutralize added H+ or OH- ions.

Acid-Base Neutralization

Acid-base neutralization is a chemical reaction in which an acid and a base react to form water and a salt. This reaction is central to the functioning of buffer solutions. When a strong acid is added to a buffer, the buffer's weak base (the conjugate base in the buffer) reacts with the strong acid's protons (\text{H+ ions}) to form water and replenish the weak acid part of the buffer.

Typical Neutralization Reaction:

The general neutralization reaction in the context of a buffer system can be represented as:
\[\ce{HA} (weak acid) + \ce{OH-} (from a strong base) \rightarrow \ce{A-} (conjugate base) + \ce{H2O} (water)\]
Such neutralization serves to maintain the pH by using up the added strong acid or base, to an extent determined by the buffer capacity. When the capacity is exceeded, the pH shall shift, indicating the buffer is no longer able to neutralize the added substances effectively.

pH Calculation

The pH calculation is a process used to determine the acidity or alkalinity of a solution. It is mathematically expressed as the negative logarithm to the base 10 of the hydrogen ion concentration:
\[\mathrm{pH} = -\log[\ce{H+}]\]
In buffer solutions, the Henderson-Hasselbalch equation provides a useful pH calculation method. It relates the pH to the pKa and the molar concentrations of the weak acid (\text{HA}) and its conjugate base (\text{A-}):
\[\mathrm{pH} = \mathrm{pKa} + \log\left(\frac{[\ce{A^-}]}{[\ce{HA}]}\right)\]
This equation helps predict the pH change that occurs when acids or bases are added to a buffer. When a strong acid is added, [H+] increases, and the ratio of [\(\ce{A^-}\)] to [\(\ce{HA}\)] changes, altering the pH in accordance with the logarithmic relationship stated in the equation.

Conjugate Acid-Base Pair

A conjugate acid-base pair consists of two species found in a solution that can transform into each other by the gain or loss of a proton (H+ ion). This concept is fundamental to understanding buffer solutions. A buffer is usually made up of a weak acid and its conjugate base. When a strong acid is added to the buffer, it reacts with the conjugate base to form the weak acid.

The weak acid (\text{HA}) and its conjugate base (\text{A-}) interact with the added H+ ions in a reversible reaction, which can be written as:
\[\ce{HA (aq) + H+ (aq) <=> A^- (aq) + H2O (l)}\]
This acid-base pair is responsible for the buffer's capability to resist changes in pH. When the balance between the conjugate acid and base is disturbed by adding more H+ ions, the equilibrium shifts to counteract this change, thereby regulating the pH within a certain range.