A solution contains 0.022 M Ag and 0.033 M Pb2 . If you add Cl–, AgCl and PbCl2 will begin to precipitate. What is the concentration of Cl– required, in molarity, when
A. AgCl precipitation begins?
B. AgCl precipitation is 99.99% complete?
C. PbCl2 precipitation begins?
D. PbCl2 precipitation is 99.99% complete?
Finally, give the concentration range of Cl– for the complete separation of Ag and Pb2 .
E. Concentration of Cl– at the start: F. Concentration of Cl– once complete
For A I got 8.18x10^-9 and for B I got 0.0227. A is wrong but I'm not sure why..? and I need help on how to approach the rest of this question please! Really confused.
2 answers
Oh and the Ksp values for AgCl and PbCl2 is 1.8x10^-10 and 1.7x10^-5, respectively.
Thanks for the Ksp values. My book had them slightly different. This way I can compare your values.
For A, I think your answer is right; you just reported too many significant figures. You're using Ksp and (Ag^+) with 2 s.f.; I suspect 8.2E-9 is what you are to report.
B. When AgCl is 99.99% that means the Ag^+ is 0.01% in solution or 0.0001 x 0.022 = ? that is left in solution. Use Ksp to calculate Cl^- at this point. I think 8.2E-5M is B.
For C I would round that to 0.023M.That should get you started.
For A, I think your answer is right; you just reported too many significant figures. You're using Ksp and (Ag^+) with 2 s.f.; I suspect 8.2E-9 is what you are to report.
B. When AgCl is 99.99% that means the Ag^+ is 0.01% in solution or 0.0001 x 0.022 = ? that is left in solution. Use Ksp to calculate Cl^- at this point. I think 8.2E-5M is B.
For C I would round that to 0.023M.That should get you started.
1 answer