The ratio is larger than allowed by the equilibrium law, and the acetate ions will combine with the hydrogen ions to form acetic acid more rapidly than the ions are formed from it, until equilibrium is reëstablished. In a molar solution of acetic acid, 0.42% of the acid is ionized, and we have [CH_{3}CO_{2}^{−}] × [H^{+}] / [CH_{3}CO_{2}H] = 0.0042 × 0.0042 / 0.9958 = 1.8E−5. If an equivalent quantity of sodium acetate should be added, ‹i.e.› one mole or 82 grams of the salt per liter, the salt being ionized to the extent of 53%, we would have 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] [CH_{3}CO_{2}^{−}] = 0.53, [H^{+}] = 0.000,034 and [CH_{3}CO_{2}H] = 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: