Since activity is a measure of the effective concentration of solute in a solution, it can be used to describe solubilities of slightly soluble solutes in ionic solutions. In other words, we can investigate the effects of ionic strength on solubility.

For example, since copper (II) sulfide is only somewhat soluble in water, we can write its dissolution as an equilibrium:

For example, since copper (II) sulfide is only somewhat soluble in water, we can write its dissolution as an equilibrium:

CuS(s) ↔ Cu

^{2+}(aq) + S^{2-}(aq)The thermodynamically correct equilibrium constant is then

K

_{sp}=a_{Cu}a_{S}=γ_{Cu}γ_{S}m_{Cu}m_{S}=γ_{±}^{2}K_{sp}^{obs}The parameter K

^{obs}is the observed (or apparent) equilibrium constant and will be the one that is measured by standard measurements like titration. Since this is a 1:1 ionic compound, K_{sp}is related to S, the molal solubility, by:K

Combining the two equations above, we can solve for the ratio of solubilities:

_{sp}= S^{2}and K_{sp}^{obs}= (S^{obs})^{2}Combining the two equations above, we can solve for the ratio of solubilities:

S/S

Therefore, for mean activity coefficients less than one, we see an increase in solubility with respect to water [salting-in] and a corresponding solubility decrease when γ

^{obs}= γ_{±}Therefore, for mean activity coefficients less than one, we see an increase in solubility with respect to water [salting-in] and a corresponding solubility decrease when γ

_{±}is greater than one [salting-out].

## 2 comments:

hi,

i may have done it wrong, but for #2 hw 2 i got 0.450 instead of 0.900. is the answer really 0.900?

nevermind. i got it.

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