the approximate balancing of two (or more) opposing factors.[303] When we leave the range of concentrations mentioned, and go to more concentrated solutions, these factors seem to be less well balanced and the validity of the principle ceases.[304] For the present it will be safe to consider the principle as an empirical one, holding for solutions of total salt content up to 0.25 or 0.3 molar.[305] For quite dilute solutions the effect of the electrolyte on the solvent would be negligible, and only to such solutions would the theoretical derivation brought forward by Washburn be applicable. «Influence of a Common Ion.»—For a saturated aqueous solution of silver acetate at a given temperature, the product of [p145] the ion concentrations may be considered a constant, [CH_{3}COO^{−}] × [Ag^{+}] = K_{S.P.}. In such an aqueous solution, containing no foreign salts, the concentration of the silver-ion is equal to the concentration of the acetate-ion, since a molecule of silver acetate, when it ionizes, gives one silver ion for every acetate ion formed. The numerical value of the solubility-product may then be calculated, if the solubility of the salt and its degree of ionization are known. For instance, at 16° one liter of water dissolves 10.07 grams of silver acetate, that is, 10.07 / 167, or 0.0603 gram-molecule (mole). Conductivity measurements show that 70.8% of the salt is ionized in such a solution, and consequently the concentration of the silver-ion is 0.708 × 0.0603, or 0.0427. The concentration of the acetate-ion is the same, and the value of the solubility-product constant, obtained by inserting these quantities in the above equation, is K_{S.P.} = 0.0427 × 0.0427 = 0.00182. Now, if, to the saturated solution of the silver acetate, there are added a few drops of a concentrated solution of sodium acetate or some crystals of solid sodium acetate, the concentration of the acetate-ion is thereby increased and the condition of equilibrium in the solution is disturbed: ‹x› [CH_{3}COO^{−}] × [Ag^{+}] > K_{S.P.}. The concentration of the acetate-ion having been increased, the ion will combine more rapidly than before with the silver-ion, and the concentration of the ‹nonionized salt› will be ‹increased›. The solution being already saturated with nonionized silver acetate, the excess formed must be ‹precipitated›. As a matter of fact, a precipitate of silver acetate is readily obtained in this way (‹exp.›). Precipitation will cease when sufficient silver acetate has crystallized out to make the product of the concentrations of the ions again equal to the solubility-product constant. If, after the crystallization is complete and equilibrium has been reëstablished, the acetate-ion is ‹x′› times as concentrated as it was in the pure aqueous solution, the concentration of the silver-ion must be reduced to 1 / ‹x′› its original value: ‹x′› [CH_{3}COO^{−}] × [Ag^{+}] / ‹x′› = K_{S.P.}.