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[c] [2] For example, a hypothetical weak acid having K a = 10 −5, the value of log K a is the exponent (−5), giving pK a = 5. For acetic acid, K a = 1.8 x 10 −5, so pK a is 4.7. A higher K a corresponds to a stronger acid (an acid that is more dissociated at equilibrium).
From the titration of protonatable group, one can read the so-called pK a 1 ⁄ 2 which is equal to the pH value where the group is half-protonated (i.e. when 50% such groups would be protonated). The pK a 1 ⁄ 2 is equal to the Henderson–Hasselbalch pK a (pK HH a) if the titration curve follows the Henderson–Hasselbalch equation. [14]
For example, acetic acid is a weak acid which has a = 1.75 x 10 −5. Its conjugate base is the acetate ion with K b = 10 −14 / K a = 5.7 x 10 −10 (from the relationship K a × K b = 10 −14 ), which certainly does not correspond to a strong base.
In chemistry and biochemistry, the Henderson–Hasselbalch equation = + ([] []) relates the pH of a chemical solution of a weak acid to the numerical value of the acid dissociation constant, K a, of acid and the ratio of the concentrations, [] [] of the acid and its conjugate base in an equilibrium.
For example, if a macromolecule M has three binding sites, K′ 1 describes a ligand being bound to any of the three binding sites. In this example, K′ 2 describes two molecules being bound and K′ 3 three molecules being bound to the macromolecule. The microscopic or individual dissociation constant describes the equilibrium of ligands ...
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