Final answer:
We stop a titration when the indicator changes color to accurately determine the equivalence point. Adding titrant beyond this point can result in an excess of titrant, leading to inaccurate pK values. Indicators are chosen to have a color change near the expected equivalence point pH for precision.
Step-by-step explanation:
During a titration, we stop adding the titrant when the indicator changes color because this signifies that we've reached the equivalence point of the reaction, where the amount of acid equals the amount of base. If we continue to add titrant beyond this point, we would have an excess of the titrant in the solution. For instance, if we added extra base, we would have unreacted base present, which could lead to a higher pH than at the true equivalence point, affecting the pK value determination inaccurately. This is why it's crucial to use an indicator that has a color change close to the expected pH at the equivalence point, providing a sharp end point for the best accuracy in our titration.
Additionally, acid-base indicators change color over a range of pH values, not at a specific value. This change is due to the relative concentrations of the different ionic and non-ionic forms of the indicator, which changes according to the pH. The color change occurs over a pH range where the ratio of these forms shifts noticeably, typically over a two pH unit range, centered around the pKin (pK) of the indicator.