Final answer:
The correct assumption regarding the molar solubility of AgCl in the presence of AgNO3 is that if the concentration of AgNO3 is significantly higher than the solubility of AgCl, the Ag+ concentration would not change significantly upon the addition of AgCl. Assumptions regarding Ksp being the same as solubility and solubility independence from AgNO3 concentration are incorrect.
Step-by-step explanation:
In the context of the question regarding AgCl solubility in the presence of AgNO₃, certain key assumptions can be made.
Firstly, assumption a stating that Ksp is the same as solubility should be dismissed, as Ksp (the solubility product constant) is a specific value for a given temperature that represents the product of the molar concentrations of the ions in a saturated solution, not the solubility itself.
Secondly, assumption c, implying that solubility of AgCl is independent of AgNO₃ concentration, is incorrect since Ag+ is a common ion in both AgCl and AgNO₃, and its presence would affect the solubility of AgCl due to the common ion effect.
Lastly, option d suggests that the concentration of Ag+ does not change significantly upon the addition of AgCl to AgNO₃ solution. This might be reasonable if the AgNO₃ concentration is much greater than the solubility of AgCl, as the change in Ag+ concentration would be minimal in comparison to the initial AgNO₃ concentration.
Understanding the solubility product, Ksp, is crucial when discussing precipitation and solubility equilibria. AgCl in pure water at 25°C has a solubility (x) of 1.33 × 10⁻⁵ M, therefore Ksp = x² = (1.33 × 10⁻⁵ M)² = 1.77× 10⁻¹⁰ for AgCl.
However, when AgCl is added to an existing AgNO₃ solution, the solubility may not be equivalent to that in pure water due to the presence of additional Ag+ ions.