What will be the solubility product constant, Q, for the precipitation reaction involving 64.5 mL of 0.045 M Ni(NO3)2 and 95.5 mL of 0.025 M Li2CO3? Ni(NO3)2 (aq) + Li2CO3 (aq) →

Give the expression for the solubility product constant for AgCl, based on the following dissolution equation:  AgCl(s)  Ag+(aq) + Cl-(aq).

A highly acidic (pH< 0.5) industrial wastewater has an initial AlP" concentration of 10 ppm. An engineer decides that strong base should be added to the wastewater in order to raise the pH so that most of the aluminum will precipitate and be removed from solution. The pertinent solubility reaction from Appendix A of our text is: Al(OH)3 (s) → Al* + 30H¯ with pK, = 32.9 a. What is the lowest pH needed to precipitate all but 50 ppb of the aluminum from the wastewater? Assume equilibrium conditions can be achieved. b. Suppose the engineer decides that the pH should be raised to 7. What equilibrium concentration of aluminum would be left in solution? Report your answer in ppb and moles per liter.

Among the solubility rules previously discussed is the statement: Carbonates, phosphates, borates, andarsenates—except those of the ammonium ion and the alkali metals—are insoluble. (a) Write the expression for the equilibrium constant for the reaction represented by the equationCaCO_3(s) ⇌ Ca^2+(aq) + CO3^1−(aq). Is Kc > 1, < 1, or ≈ 1? Explain your answer. (b) Write the expression for the equilibrium constant for the reaction represented by the equation3Ba^2+(aq) + 2PO_4^3−(aq) ⇌ Ba_3(PO_4)_2(s). Is Kc > 1, < 1, or ≈ 1? Explain your answer.

At 25 °C, only 0.0110 mol of the generic salt AB2 is soluble in 1.00 L of water. What is the Ksp of the salt at 25 °C? AB, (s) = A2+(aq) + 2B¯(aq)

A student prepares a solution for titration by adding a 10.00 mL aliquot of a solution saturated with Ca(OH)2 at 82.9°C to 25.0 mL of deionized water.  This solution is titrated with a standardized 0.01252 M HCl solution.  The equivalence point is observed after 17.94 mL of the HCl solution has been added. Calculate the moles and molarity of OH-  and Ca2+  in the 10.00ml sample.

a A volume of 500.0 mL of 0.130 M NaOH is added to 585 mL of 0.250 M weak acid (K₂ = 6.46 x 10-5). What is the pH of the resulting buffer? HA(aq) + OH¯(aq) → H₂O(1) + A¯(aq) pH =

7.  The separate samples of nitric and acetic acids are both titrated with a 0.100 M solution of NaOH(aq). (Y) 25.0mL of 1.0 M HNO3(aq)                      (Z) 25.0 mL of 1.0 M CH3COOH(aq)   Determine whether each of the following statements concerning this titration is true or false. (A) A larger volume of NaOH(aq) is needed to reach the equivalence point in the titration of HNO3 (B) The pH at the equivalence point in the HNO3 titration will be lower than the pH at the equivalence point in the CH3COOH titration (C) Phenolphthalein would be a suitable indicator for both titrations Group of answer choices a.) A) False B) True C) True b.) A) False B) False C) True c.) A) False B) True C) False d.) A) True B)True C)True

Write a balanced net ionic equation to show why the solubility of Mg(OH)2(s) increases in the presence of a strong acid and calculate the equilibrium constant for the reaction of this sparingly soluble salt with acid.Be sure to specify states such as (aq) or (s). Mg(OH)2(s) + 2HCl(aq)   MgCl2(aq) + 2H2O(l) K =

The solubility of iron(II) hydroxide, Fe(OH)2, is 1.43×10–3 gram per liter at 25 °C.  (a) Write a balanced equation for the solubility equilibrium.   (b) Write the expression for the solubility product constant, Ksp, and calculate its value.   (c) Calculate the pH of a saturated solution of Fe(OH)2 at 25 °C.