15.2:
Acid/Base Strengths and Dissociation Constants
Acids and bases can be categorized by whether they are a strong acid, a strong base, a weak acid, or a weak base. There are very few strong acids and bases so the majority of acids and bases are weak.
A strong acid, like hydrochloric acid, completely dissociates into hydrogen ions and chloride ions when dissolved in water. A strong base, like sodium hydroxide, dissociates completely into sodium ions and hydroxide ions.
Weak acids and bases partially dissociate and are present in both ionized and un-ionized forms. For example, both acetic acid and its weak conjugate base, acetate, are found in an aqueous solution.
The degree of dissociation of a weak acid or base can be measured using its equilibrium constant. The equilibrium constant for weak acids has a special name, the acid dissociation constant, or, Ka.
For a generic weak acid HA, Ka at a given temperature can be calculated by an equilibrium equation: dividing the concentration of products, A ion and hydronium ion, by the concentration of reactants, HA and water.
As water is liquid and its concentration remains nearly unchanged in the reaction, it is excluded from the equation.
The higher the Ka, the stronger the acid. Nitrous acid is stronger than acetic acid because the Ka of nitrous acid is larger than the Ka of acetic acid.
The equilibrium constant for weak bases, the base dissociation constant, or Kb, acts in a similar manner to Ka. For a generic weak base B, Kb at a given temperature can be determined by dividing the concentration of products, BH ion, and hydroxide ion, by the concentration of reactant, B.
Like acids, the strength of the bases is also directly proportional to the Kb. For example, ethylamine is relatively stronger than urea because the Kb of ethylamine is larger than the Kb of urea.
The strength of an acid can also be expressed in terms of percent ionization. The percent ionization of an acid can be calculated by dividing hydronium ion concentration at equilibrium by the initial acid concentration and multiplying it by a hundred.
Similarly, the percent ionization for bases can be calculated by dividing the hydroxide ion concentration at equilibrium by the initial concentration of base and multiplying it by a hundred. The higher the percent ionization, the stronger the acid or base.
15.2:
Acid/Base Strengths and Dissociation Constants
The relative strength of an acid or base is the extent to which it ionizes when dissolved in water. If the ionization reaction is essentially complete, the acid or base is termed strong; if relatively little ionization occurs, the acid or base is weak. There are many more weak acids and bases than strong ones. The most common strong acids and bases are listed below:
| Strong Acids | Strong Bases |
| HClO4 | LiOH |
| HCl | NaOH |
| HBr | KOH |
| HI | Ca(OH)2 |
| HNO3 | Sr(OH)2 |
| H2SO4 | Ba(OH)2 |
The relative strengths of acids may be quantified by measuring their equilibrium constants in aqueous solutions. In solutions of the same concentration, stronger acids ionize to a greater extent and so yield higher concentrations of hydronium ions than do weaker acids. The equilibrium constant for an acid is called the acid-ionization constant, Ka. For the reaction of an acid HA:
the acid ionization constant is written as
where the concentrations are those at equilibrium. Although water is a reactant in the reaction, it is the solvent as well, so we do not include [H2O] in the equation. The larger the Ka of an acid, the larger the concentration of H3O+ and A− relative to the concentration of the nonionized acid, HA, in an equilibrium mixture, and the stronger the acid. An acid is classified as “strong” when it undergoes complete ionization, in which case the concentration of HA is zero and the acid ionization constant is immeasurably large (Ka ≈ ∞). Acids that are partially ionized are called “weak,” and their acid ionization constants may be experimentally measured.
To illustrate this idea, three acid ionization equations and Ka values are shown below. The ionization constants increase from first to last of the listed equations, indicating the relative acid strength increases in the order CH3CO2H < HNO2 < HSO4−.
Another measure of the strength of an acid is its percent ionization. The percent ionization of a weak acid is defined in terms of the composition of an equilibrium mixture:
where the numerator is equivalent to the concentration of the acid's conjugate base (per stoichiometry, [A−] = [H3O+]). Unlike the Ka value, the percent ionization of a weak acid varies with the initial concentration of acid, typically decreasing as concentration increases.
Just as for acids, the relative strength of a base is reflected in the magnitude of its base-ionization constant (Kb) in aqueous solutions. In solutions of the same concentration, stronger bases ionize to a greater extent, and so yield higher hydroxide ion concentrations than do weaker bases. A stronger base has a larger ionization constant than does a weaker base. For the reaction of a base, B:
the ionization constant is written as
The inspection of the data for three weak bases presented below shows the base strength increases in the order NO2− < CH2 CO2− < NH3.
As for acids, the relative strength of a base is also reflected in its percent ionization, computed as
but will vary depending on the base ionization constant and the initial concentration of the solution.
This text is adapted from Openstax, Chemistry 2e, Section 14.3: Relative Strengths of Acids and Bases.