Question

1. (15 pts) Use Ksp value (1E-32) and solubility reaction for aluminum hydroxide, AI(OH)3, found in Table 2.3 (p.62). pH is fixed at 6.0. What is the solubility of aluminum hydroxide? Report your result as the concentration of the aluminum metal cation in M units. Assume an infinite solid. Explain briefly how your result does or does not match the curves shown in the figure in class notes (see the graph for amorphous aluminum hydroxide solubility as a function of pH) Aluminum hydroxide is dissolved in pure water. At equilibrium, what is the metal ion concentration (in M) and the pH? Some aquatic species can be killed by aluminum ion concentrations as low as 1 mg/L. Is there a pH range indicated by the amorphous aluminum hydroxide solubility graph that could achieve dissolve aluminum ions this low? If so, what is the range?

Answer 1

K_sp Al(OH)_3 = 1 times 10^-32 (Graph is not given in image) Al(OH)_3 rightleftharpoons AL^3+ + 3 O H^- P^H = 6.0 Rightarrow POH = 8 [OH^-] = 10^-8 M K_sp = [AL^3+] [O H^-]^3 1 times 10^-32 = [AL^3+] [10^-8]^3 [AL^3+] = 10^-8 M In pure water [H^+] = [O H^-] = 10^-7 M K_sp = [AL^3+] [10^-7]^3 1 times 10^-32 = [AL^3+] [10^-21] [AL^3+] = 1 times 10^-11 M max [AL^3+] = 1 mg/L [AL^3+] = 1 times 10^-3/27 = 3.7 times 10^-5 mol/ltr K_sp = [AL^3+] [O H^-]^3 1 times 10^-32 = (3.7 times 10^-5) [S]^3 (S)^3 = 2.7 times 10^-28 S = 6.47 times 10^-10 M [H^+] = 6.47 times 10^-10 M P_H = -log [H^+] = 9.19 So P^H < 9.19 to keep AL^3+ under control