Final answer:
Exceeding the titration endpoint of a saturated Ca(OH)2 solution results in a reported Ksp that is too low, due to the disruption of equilibrium and precipitation of additional Ca(OH)2, lowering the ionic concentrations.
Step-by-step explanation:
If the endpoint in the titration of a saturated Ca(OH)₂ solution with a standardized HCl solution is surpassed, the reported Ksp of Ca(OH)₂ will be reported too low. This occurs because once the endpoint is exceeded, additional HCl reacts with the saturated solution, affecting the equilibrium. According to Le Chatelier's principle, the equilibrium will shift to accommodate the excess HCl by precipitating more Ca(OH)₂, thus decreasing the concentrations of Ca²⁺ and OH⁻⁺ ions in the solution.
This results in a lower ionic product (Q) than the actual solubility product constant (Ksp) because the reaction has been driven to the left, creating a supersaturated solution. This means that more solid is present than would be in a saturated solution at equilibrium, which leads to an underestimation of the Ksp when calculated based on the concentrations after surpassing the endpoint.