You note that the Ksp of Ba3(PO4)2 is much smaller than that of Ba(OH)2. Therefore, you wonder if adding a phosphate species, such as H3PO4, can precipitate out more of the barium ions (Ba²+) from the solution (and then you could filter out the solid to have a smaller waste volume!). Therefore, you look up another equilibrium constant, shown for equation below: H3PO4 (aq)=3 H+ (aq) + PO¾-¯ (aq) K₁ = 2.0 x 10-22 15. If you were to add H3PO4 to your waste solution to attempt to precipitate Ba3(PO4)2, what would be the equilibrium constant for the net reaction (shown below)? 3 Ba(OH)2 (s) + 2 H3PO4 (aq) = 6 H2O (l) + Ba3(PO4)2 (5)

What is the solubility product expression for Co(OH)3? Ksp= [Co3+][30H] Ksp= 3[Co3+][OH73 Ksp= [Co3+][OH73 Ksp= [Co3+][30H]3 Ksp= [Co3*][OH]

The equilibrium constant for the formation of Cu(CN)42- is 2.0 x 1030. Calculate the value of pCu2+, or -log[Cu2+], if 2.21 g of CuCl2 is dissolved in 1.000 L of a 0.926 M solution by NaCN. The addition of CuCl2 does not affect the volume (the final volume is always 1,000 L).

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.