Question

1. How is the molar solubility of a slightly soluble salt affected by the addition of an ion common to the salt equilibrium? 2. A 3.11 mL volume of a standardized 0.0025 M HCl solution titrated 25.0 mL of a saturated Mg(OH)2 solution to the methyl orange endpoint. Calculate the Ksp of Mg(OH)2. 3. If the endpoint in the titration of a saturated Ca(OH)2 solution with a standardized HCI solution is surpassed, will the reported Kap of Ca(OH)2 be reported too high or too low? Explain 4. The Ksp of MgF is 6.4 x 109. What is the molar solubility of MgF in a solution of o.10 M MgCl2?. 5. A water droplet clings to the inner wall of a dirty buret between the initial and final volumes in the titration of a saturated Ca(OH)2 solution. Will the molar solubility of the Ca(OH)2 be reported too high, too low, or be unaffected by this technique error?

Answer 1

1)

Adding a ion ( cation or anion) common to sparingly soluble salt will shifts the solubility equilibrium towards backward direction which results in decrease of solubility of sparingly soluble salt. This is called COMMON ION EFFECT.

COMMON ION EFFECT - - - - > It states that there is a shift in the equilibrium according to Le chatelier when there is an addition of a ion already present in equilibrium reaction, which leads to decrease in the solubility and salt will finally precipitated out.