From this experimental result, the equilibrium constant, which is called the ‹ionization constant› of the acid, is found to have the value 18.2E−6. If the ratio [H^{+}] × [CH_{3}CO_{2}^{−}] / [CH_{3}CO_{2}H] really is a constant, the same value, within the limits of experimental errors, should be obtained from acetic acid in other concentrations. Now, if the above solution is diluted to ten times its volume, the concentration of the acid is made 0.01 mole ‹per liter›, the conductivity [p099] is found to have increased to 14.5 reciprocal ohms, and the percentage of ionized acid is then 14.5 / 347, or 4.17. Here, [H^{+}] and [CH_{3}CO_{2}^{−}] = 0.01 × 0.0417 and [CH_{3}CO_{2}H] = 0.01 × 0.9583. Inserting these values in our general equation and calculating the result, we obtain 18.1E−6 as the value of the constant. In the following table[180] are given the molar conductivities, Λ (column 2), of acetic acid of varying concentrations, ‹m› (column 1). The degrees of ionization, α, and the ionization constant, calculated according to the equilibrium equation, are given in columns 3 and 4. IONIZATION OF ACETIC ACID. Λ_{∞} = 347. m. Λ. 100 α. K. 0.1 4.67 1.34 18.2E−6 0.08 5.22 1.50 18.3E−6 0.03 8.50 2.45 18.5E−6 0.01 14.50 4.17 18.1E−6 It is evident, that a constant value is found for the ratio [H^{+}] × [CH_{3}CO_{2}^{−}] / [CH_{3}CO_{2}H] and that the ionization of acetic acid, in these dilute solutions, obeys the law of chemical equilibrium.[181] The equilibrium constant expresses in definite, quantitative terms the tendency of acetic acid to ionize in dilute solution. Examination of other acids shows that there is an enormous range in the values found for their respective ionization constants. The constants are the best ‹measure› of the ‹strength› of the ‹acids› as acids. Obviously, the more readily acids in equivalent solutions ionize, the greater will be the concentration of the hydrogen-ion to which the characteristic acid properties are due, and the more pronounced (stronger) will be the exhibition of these properties. From the ionization constants one may calculate, for instance, the proportion in which two competing acids will neutralize a base, when the latter is used in quantity insufficient to neutralize both acids. [p100] Inspection of the equation for acetic acid, which is the typical equilibrium equation for all ‹monobasic› acids, shows that the greater the degrees of ionization of acids are in equivalent solutions, ‹i.e.› the greater the concentrations of the hydrogen-ion which their ionization produces in equivalent solutions, the larger will be the values of their ionization constants. The acids with the ‹larger› constants are, then, the ‹stronger› acids.