If 2.5mL of 0.30M AgNO_3 is mixed with 7.5mL of 0.015M Na_2SO_4, should a precipitate of Ag_2SO_4 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^(+) + SO_4^(2-) rightleftharpoonsAg_2SO_4(s)darr

Q_"the ion product"=[Ag^+]^2[SO_4^(2-)]

And now, we must determine the individual concentrations:

[Ag^+] = (2.5xx10^(-3)*Lxx0.30*mol*L^-1)/((2.5+7.5)xx10^-3L) = 7.50xx10^-2*mol*L^-1.

[SO_4^(2-)] = (7.5xx10^(-3)*Lxx0.015*mol*L^-1)/((2.5+7.5)xx10^-3L) = 1.13xx10^-2*mol*L^-1.

Q_"the ion product"=(7.50xx10^-2)^(2)(1.13xx10^-2) = 6.4xx10^-5.

Since Q_"the ion product"=6.4xx10^-5>K_"sp"=1.2xx10^-5, precipitation of silver sulfate should occur until equilibrium is satisfied, and Q=K_"sp"

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 K_"sp". At higher temperatures, how would you expect K_"sp" to evolve? Would solubility of silver sulfate increase or decrease?

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