Suppose that a buffer contains equal amounts of a weak acid and its conjugate base. What happens to the relative amounts of the weak acid and conjugate base when a small amount of strong acid is added to the buffer? What happens when a small amount of strong base is added?

Understanding the relationship between a weak acid and its conjugate base is essential for grasping buffer solution chemistry. A weak acid only partially dissociates in water, yielding a mixture of undissociated acid molecules and their corresponding ions.

A weak acid, HA, can be represented by the equilibrium equation: \[ HA \rightleftharpoons H^+ + A^- \].
The ion A- is known as the conjugate base of the weak acid. It can accept a proton (H+) to form the acid, thus playing a crucial role in the buffering action. This acid-base pair can neutralize added strong acids or bases to the solution, which minimizes changes in pH. The ability of these pairs to sustain a stable pH is why they are commonly used to create buffer solutions.

Buffer capacity refers to the amount of strong acid or strong base a buffer solution can absorb without undergoing a significant change in pH. It is influenced by the concentration of the acid/base pair; higher concentrations provide a larger capacity to neutralize added acids or bases.

The buffer capacity is also related to the proximity of the pH to the pKa of the weak acid. Maximum buffering capacity is achieved when the pH is equal to the pKa, where the concentrations of the weak acid and conjugate base are equal. This balance is delicate, and thus, if a buffer is overwhelmed with too much strong acid or base, it can exceed its capacity, leading to a more drastic shift in pH.

When you introduce a strong acid to a buffer solution, the high concentration of H+ ions from the strong acid reacts with the conjugate base. This neutralization reaction forms more of the weak acid, effectively absorbing the added H+ ions and minimizing the change in pH. The chemical reaction can be represented as: \[ A^- + H^+ \rightarrow HA \].
Despite the transformation of some conjugate base into the weak acid, if the buffer capacity is adequate, the pH alteration will be minuscule. This is due to the equilibrium between the weak acid and its conjugate base, which still persists despite some shift in their relative amounts. This dynamic is a perfect illustration of a buffer's resistance to pH changes.

Conversely, the addition of a strong base to a buffer involves the reaction of OH- ions with the weak acid component of the buffer. This interaction produces water and additional conjugate base, as depicted in the reaction: \[ HA + OH^- \rightarrow A^- + H_2O \].
The buffer solution responds by using up some of the weak acid to neutralize the OH- ions, yielding more conjugate base but leaving the overall pH relatively unaffected, assuming the buffer's capacity is not surpassed. It is this characteristic that makes buffers an invaluable tool in maintaining pH stability in a variety of chemical and biological systems.