| (0.1 × 0.0134)2 | = Kionization = 18.2E−6. |
| 0.1 × 0.9866 |
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+] × [CH3CO2−] / [CH3CO2H] 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 [CH3CO2−] = 0.01 × 0.0417 and [CH3CO2H] = 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+] × [CH3CO2−] / [CH3CO2H] 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.
The Ionization of Various Acids.
In the first place, for the strongest acids, such as hydrochloric, nitric, hydrobromic and similar acids, chemists have been unable to determine ionization constants on the basis of the law of chemical equilibrium. Strong acids, strong bases and most salts (see pp. [106]–[8], below), the three classes comprising all the very readily ionizable electrolytes, do not give constants when the values of the equilibrium ratio,[182] [Cation] × [Anion] / [Molecules], are calculated for different concentrations, and they therefore do not ionize simply in accordance with the law of chemical equilibrium. The reasons for this abnormal behavior will be discussed presently (p. [108]), when other necessary facts are before us. In order to have, at least, a rough basis for comparison of these strong acids with the weak ones, which do obey the law of chemical equilibrium, the table will give for the strong acids the value of the above ratio as calculated from their ionization in 0.1 molar solutions.
The Ionization of Polybasic Acids.
[H+]2 × [CO32−] / [H2CO3] = K.