If 2.5mL of 0.30M AgNO3 is mixed with 7.5mL of 0.015M Na2SO4, should a precipitate of Ag2SO4 form? (Ksp = 1.2x10−5)

1 Answer
anor277
Jul 3, 2016

Quite probably.

Explanation:

We need to work out the ion product for the following reaction, under the given conditions:

2Ag++SO2−4⇌Ag2SO4(s)⏐⏐↓

Qthe ion product=[Ag+]2[SO2−4]

And now, we must determine the individual concentrations:

[Ag+] = 2.5×10−3⋅L×0.30⋅mol⋅L−1(2.5+7.5)×10−3L = 7.50×10−2⋅mol⋅L−1.

[SO2−4] = 7.5×10−3⋅L×0.015⋅mol⋅L−1(2.5+7.5)×10−3L = 1.13×10−2⋅mol⋅L−1.

Qthe ion product=(7.50×10−2)2(1.13×10−2)=6.4×10−5.

Since Qthe ion product=6.4×10−5>Ksp=1.2×10−5, precipitation of silver sulfate should occur until equilibrium is satisfied, and Q=Ksp

Please don't trust my arithmetic. The approach I took (I think) was sound; we had to add volumes and recalculate concentrations.

Temperature was not referred to in this question; we would assume room temperature for Ksp. At higher temperatures, how would you expect Ksp to evolve? Would solubility of silver sulfate increase or decrease?

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