Adding a buffer solution to an acid-base reaction between a weak acid and its conjugate base will not change the equilibrium constant K_a of the reaction. However, it will help to maintain the pH of the solution by resisting changes in the concentration of H+ and OH- ions. Let's consider an example to illustrate this.Suppose we have a weak acid, HA, and its conjugate base, A-. The equilibrium reaction can be represented as:HA <=> H+ + A-The equilibrium constant K_a for this reaction is given by:K_a = [H+][A-] / [HA]Now, let's say we add a buffer solution containing the same weak acid HA and its conjugate base A- . The buffer solution will have a certain concentration of HA and A-, which we can represent as [HA]_buffer and [A-]_buffer.Step 1: Calculate the initial pH of the solutionTo find the initial pH of the solution, we can use the Henderson-Hasselbalch equation:pH = pK_a + log [A-] / [HA] where pK_a is the negative logarithm of the equilibrium constant K_a.Step 2: Determine the change in equilibrium positionWhen the buffer solution is added, the concentrations of HA and A- will change. The new concentrations can be represented as:[HA]_new = [HA] + [HA]_buffer[A-]_new = [A-] + [A-]_bufferStep 3: Calculate the final pH of the solutionNow, we can use the new concentrations of HA and A- to find the final pH of the solution using the Henderson-Hasselbalch equation:pH_final = pK_a + log [A-]_new / [HA]_new This final pH value will be close to the initial pH value due to the buffering effect of the added buffer solution. The equilibrium constant K_a remains unchanged throughout the process, as it only depends on the nature of the weak acid and its conjugate base, not on the concentrations of the species involved in the reaction.