The Elements of Qualitative Chemical Analysis, vol. 1, parts 1 and 2. / With Special Consideration of the Application of the Laws of Equilibrium and of the Modern Theories of Solution.

by Julius Stieglitz

0.534 × 0.0042 / 0.9958 > 1.8E−5.

The equilibrium constant will be satisfied (and the velocity of ionization of the acetic acid will become equal to the velocity of its [p113] formation from its ions) when[217] [CH₃CO₂⁻] = 0.53, [H⁺] = 0.000,034 and [CH₃CO₂H] = 0.999,966. If we use two characteristic places in the decimals, we have

(0.53 × 0.000034) / 1 = 1.8E−5.

The most significant fact in the new condition of equilibrium is the extremely small concentration of hydrogen-ion in the solution. Since the acid properties of acetic acid are due to its forming hydrogen-ion, we would conclude that such properties of acetic acid are very much weakened by the presence of its own salts. This conclusion has been fully verified by careful quantitative measurements and can be demonstrated as follows[218]:

Exp. To two of three portions of a dilute solution of methyl orange equal quantities of acetic acid are added (e.g. 0.5 c.c. molar acid), the third portion being reserved to show the color of the neutral solution of the indicator. Now, if into one of the solutions, to which acetic acid has been added, a few crystals of sodium acetate are gradually dropped, the color reverts gradually to the color of the original neutral solution; the concentration of hydrogen-ion becomes so small, that it does not visibly affect this indicator, which is not very sensitive to acids, (H⁺).[219]

The solution, according to the views expressed, should still be very slightly acid, and by the use of an indicator (litmus paper) which is much more sensitive to the hydrogen-ion than is methyl orange, no difficulty is found in recognizing this fact also. As a matter of experiment, then, an acid like acetic acid is very much weaker in the presence of its own salts than in their absence, and the equilibrium constant and the concentrations of the components used determine the extent to which the hydrogen-ion is suppressed. The same must be true for all weak acids and similar relations must [p114] hold for all weak bases[220]—in general, weak acids and weak bases are very much weakened by the addition of their own salts.

The importance of recognizing such changes, in considering analytical reactions, may be illustrated as follows:

Exp. A solution of ferrous acetate (ferrous chloride with an equivalent quantity of sodium acetate) is treated with hydrogen sulphide; a precipitate of black ferrous sulphide is formed. A second portion of the solution is first decidedly acidified with acetic acid: hydrogen sulphide does not precipitate any ferrous sulphide. Some crystals of sodium or ammonium acetate are added to the mixture and a black precipitate of iron sulphide immediately appears around the salt as it dissolves, and on mixing the contents a heavy precipitate of the sulphide throughout the vessel is formed.

It is evident that the addition of a neutral salt, containing the same negative ion as the added acid, may completely reverse the net result of a test with hydrogen sulphide.

(3) If the concentration of the hydrogen ions in the solution of acetic acid is increased by the addition of hydrochloric, or some other strong acid, equilibrium between the acetic acid and its ions will likewise be disturbed and the new condition of equilibrium will show a suppression of the acetate-ion. Instances of such action of a strong acid in suppressing the characteristic ions of weaker acids will be discussed in detail in connection with the analytical applications of hydrogen sulphide (Chap. XI), where the action is of peculiar importance. Strong bases have a similar effect on weak bases.