During your titration, you overshot your endpoint and your solution was too pink. Would this lead to a final pKa that was too high or too low? Explain why using chemical reasoning.

When an acid-base titration is performed, the endpoint is the point at which the reaction between the acid and base is complete. In an ideal situation, the endpoint should be achieved exactly at the stoichiometric equivalence point, where the moles of acid and base are present in exact proportion. However, in reality, it is common to overshoot the endpoint slightly, leading to an excess of one of the reagents.

In the scenario you described, where the solution turned too pink, it means that you added excess base beyond the endpoint. This excess base will then react with the conjugate acid, forming a new species that is either a weak acid or a weak base, depending on the specific acid-base system involved.

In general, when a solution is too pink, with a higher concentration of the conjugate acid, it suggests the presence of a higher concentration of the weak acid species. This is because the excess base reacted with the conjugate acid to form the weak acid species. Since weak acids have a lower tendency to donate protons (H+ ions), their equilibrium constant, known as the acid dissociation constant (Ka), is smaller compared to strong acids.

Now, for the pKa value, it is defined as the negative logarithm of the Ka value. Therefore, if the weak acid concentration increases, the Ka value decreases, and consequently, the pKa value increases as well.

Therefore, in the scenario you described, when the solution turned too pink, it indicates the presence of excess weak acid species. As a result, the final pKa will be higher than expected because the concentration of the weak acid is higher due to the overshooting of the endpoint and the subsequent reaction with excess base.