Answer:
To predict if a precipitate will form when 2 solutions are mixed:
• Calculate the concentration of individual ions in the combined solution (remember to divide mol of ion by the TOTAL
solution volume)
• Calculate the ion product, Q (same form as Ksp, but ion concentrations here are not necessarily equilibrium ion
concentrations)
• Compare Q to Ksp
• If Q > Ksp, the solution is supersaturated, ion concentrations are greater than equilibrium concentrations, reaction will
proceed in reverse to reach equilibrium, precipitation will occur.
• If Q < Ksp, the solution is unsaturated, ion concentrations are less than equilibrium concentrations, reaction will proceed
forward to reach equilibrium, more solid will dissolve.
• If Q = Ksp, the solution is saturated, the solution is at equilibrium, ion concentrations are equilibrium concentrations, no
more solid will dissolve or precipitate.
Explanation:
• What minimum concentration of Na2CO3 (aq) is required to cause precipitation of BaCO3 (s) from a solution of 1.0 x 10—5
M BaCl2 (aq)?
BaCO3 (s) ↔ Ba2+ (aq) + CO3
2— (aq); Ksp = 2.6 x 10—9
logic: At the point where Q = Ksp the solution is at equilibrium. Calculate the concentration of CO3
2— that will satisfy this
relationship:
Q = [Ba2+][CO3
2—]; 2.6 x 10—9 = 1.0 x 10—5[CO3
2—]
[CO3
2—] = 2.6 x 10—4 M
This is the concentration of Na2CO3 that will result in a saturated solution (solid and ions in equilibrium). SO . . . any
[Na2CO3] greater than 2.6 x 10—4 M will result in Q > Ksp and precipitation of BaCO3.
