([CH3COO−] / y′) × y′ [Ag+] = KS.P..
The following table shows the relations when sodium acetate is added to the saturated solution of silver acetate. Column 1 gives the concentration of the sodium acetate in the solution saturated with silver acetate, column 2 the percentage of the sodium acetate that is ionized, column 3 the total concentration of silver acetate in the saturated solution, column 4 the percentage of it which is ionized, column 5 the concentration of the acetate-ion, column 6 the concentration of the silver-ion and column 7 the value of the solubility-product.
| 1 Na-Acet. | 2 100 p. | 3 Ag-Acet. | 4 100 p′. | 5 [CH3COO−]. | 6 [Ag+]. | 7 KS.P.. |
|---|---|---|---|---|---|---|
| ... | 0.0603 | 70.8 | 0.0427 | 0.0427 | 0.00182 | |
| 0.061 | 78.6 | 0.0392 | 64.5 | 0.0735 | 0.0258 | 0.00185 |
| 0.119 | 75.8 | 0.028 | 59.7 | 0.1065 | 0.0167 | 0.00179 |
| 0.239 | 70.8 | 0.0208 | 52.3 | 0.1727 | 0.0109 | 0.00188 |
The second table shows the relations when an excess of the silver-ion is present, silver nitrate having been added to the saturated silver acetate solution. The columns have the same significance as in the first table, excepting that the first column gives the concentration of silver nitrate present and the second column its degree of ionization.[306]
It is clear from these results that a difficultly soluble salt is rendered less soluble (see column 3 of the tables) by the presence of another salt, when the [p147] latter has an ion in common with the former. This conclusion has been well established[307] for a considerable number of salts.[308]
| 1 AgNO2. | 2 100 p. | 3 Ag-Acet. | 4 100 p′. | 5 [CH3COO−]. | 6 [Ag+]. | 7 KS.P.. |
|---|---|---|---|---|---|---|
| 0. | ... | 0.0603 | 70.8 | 0.0427 | 0.0427 | 0.00182 |
| 0.061 | 82.0 | 0.0417 | 64.0 | 0.0267 | 0.0767 | 0.00204 |
| 0.119 | 78.4 | 0.0341 | 58.6 | 0.0200 | 0.1142 | 0.00227 |
| 0.239 | 74.0 | 0.0195 | 51.7 | 0.0100 | 0.1809 | 0.00182 |
Applications in Analysis.
In passing, we may ask what the approximate loss of dissolved nonionized barium sulphate would amount to. The value of the ratio ([Ba2+] × [SO42−]) : [BaSO4], representing the ionization of barium sulphate, is unknown for the extreme dilution represented [p148] by the saturated solution. If we assume it to be roughly of the order 2000 : 1,[309] the solubility of nonionized barium sulphate at 18° would be roughly 0.05 milligram per liter.
As a rule, then, in the absence of complicating conditions,[310] an excess of the precipitant promotes the complete precipitation of an ionogen.