Let the solubilities of `AgCl` in `H_(2)O`, and in `0.01M CaCl_(2), 0.01M NaCl`, and `0.05M AgNO_(3)` be `S_(1),S_(2),S_(3),S_(4)`, respectively. What is the correct relationship between these quantities.

Options

A. `S_(1) gtS_(2) gt S_(3) gtS_(4)`
B. `S_(1) gt S_(2) =S_(3) gt S_(4)`
C. `S_(1) gtS_(3) gt S_(2) gt S_(4)`
D. `S_(4) gt S_(2) gtS_(3) gt S_(1)`

Answer

C

Explanation

The equation representing ionisation of `AgCl` is
`AgCl hArr Ag^(+) +Cl^(-)`
Any of the common ion `Ag^(+)` or `Cl^(-)` will tend to push the equilibrium in backward direction and thus, decrease the solubility of `AgCl`. Larger the concentrations of common ion, lesser will be the solubility.
Now, in `H_(2) O ...........................` no common ion
in `CaCl_(2) (0.01 M ) .................. [CL^(-)]=0.02 M`
in `NaCl(0.01M) ................ [Cl^(-)]=0.1`
in `AgNO_(3)(0.05) ............[Ag^(+)]=0.05`
Thus, solubility of `AgCl` should be maximum in water `(S_(1))` and it should decrease in `NaCl(S_(3))` to `CaCl_(2)(S_(2))` to `AgNO_(3)(S_(4))`.