Chemistry question about solubility products and #K_(sp)#?

A solution contains #1.0 * 10^-4# M #Pb^(2+)# and #2.0 * 10^(-3)# M #Sr^(2+)#. If a source of #SO_4^(2-)# is added to this solution, will #PbSO_4# (#K_(sp)# = #1.8 * 10^(-8)#) or #SrSO_4# (#K_(sp)# = #3.4 * 10^(-7)#) precipitate first? Specify the concentration of #SO_4^(2-)# necessary to begin precipitation of each salt.

1 Answer
Dave
Mar 19, 2018

See Below

Explanation:

Oh, #K_sp#

We just treat them individually to start. I'm going to start with #Pb^(+2)#
#PbSO_4 = Pb^(+2) + SO_4^(-2)#
I - +0.0001 0
C +s +s
E (0.0001+s) +s

#1.8xx10^(-8) = (1xx10^(-4) + s)xx(s)#
Quadratic Equation: s = 0.0000932M #SO_4^(-2)#
This tells you the concentration of #SO_4^(-2)# that will cause the #Pb^(+2)# ions to start to precipitate.

#Sr^(+2)#
#SrSO_4 = Sr^(+2) + SO_4^(-2)#
I - +0.002 0
C +s +s
E (0.002+s) +s

#3.4xx10^(-7) = (2xx10^(-3) + s)xx(s)#
Quadratic Equation: s = 0.000159M #SO_4^(-2)#
This tells you the concentration of #SO_4^(-2)# that will cause the #Sr^(+2)# ions to start to precipitate.

This means #Pb^(+2)# ions will precipitate first.

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