A pH controller that feeds NaOH to a chromium precipitation tank is set to pH 6.2, which enables most of the dissolved chromium to be removed from a metal plating waste stream and be recovered as a chromium hydroxide precipitate. If the pKsp of Cr(OH); is 30.2, what will be the equilibrium dissolved chromium concentration (in ppb) remaining in solution after the precipitate has completely settled from the treated water?

Given the following equilibria, calculate the concentration of each copper species in a solution saturated with Cu(OH)2(s) and containing [OH-] at a fixed concentration of 3.2 x 10^-7 M.   Cu(OH)2(s) Ksp= 4.8 x 10^-20   Cu(OH)^+ B1= 3.16 x 10^6   Cu(OH)2(aq) B2= 6.31 x 10^11   Cu(OH)3 ^- B3= 3.16 x 10^14   Cu(OH)4 ^2- B4= 3.98 x 10^15

Calculate the value of the equilibrium constant, K., for the reaction AgBr(s) + 2 NH, (aq) =[Ag(NH;),]* (aq) + Br (aq) K. = ? The solubility product constant, Kgp, for AgBr is 5.35 x 10-13, and the formation constant, K¢ (B2), for [Ag(NH,),I* is 1.6 x 107. K. =

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.