Last updated: 2026-06-16

The acid dissociation constant $K_a$ measures how completely an acid ionizes in water, but its values span dozens of orders of magnitude, from around $10^{2}$ for strong acids to $10^{-50}$ for the weakest. The $pK_a$ scale compresses that range into manageable numbers, just as pH does for hydrogen-ion concentration. This calculator converts $K_a$ to $pK_a$ and reports the conjugate base's $K_b$ and $pK_b$, the quantities used to rank and compare acids throughout chemistry and biochemistry.

The logarithmic definition

The $pK_a$ is the negative base-10 logarithm of the dissociation constant:

$pK_a = -\log_{10}(K_a)$

Because the relationship is logarithmic, a difference of one $pK_a$ unit corresponds to a tenfold difference in $K_a$. A lower $pK_a$ marks a stronger acid.

For the conjugate base, the dissociation constants are tied together by the ion product of water, $K_w = 1.0\times10^{-14}$ at 25 °C:

$K_b = \frac{K_w}{K_a}, \qquad pK_a + pK_b = 14$

Worked example

Acetic acid has $K_a = 1.8\times10^{-5}$:

The moderately small $K_a$ marks acetic acid as a weak acid, and its conjugate base, the acetate ion, is correspondingly a very weak base.

Why the scale is so useful

Logarithms turn awkward exponential quantities into a linear scale that is easy to compare and to use in equations. The $pK_a$ appears directly in the buffer relationship through the Henderson-Hasselbalch Calculator, where a buffer is most effective when the pH is near the acid's $pK_a$. To find the pH of a solution from its hydrogen-ion concentration, use the pH Calculator, and to see what fraction of a weak acid actually dissociates, the Percent Ionization Calculator.