The concepts of $K_a$ and $pK_a$ are closely related, yet they convey different information regarding the acidity of a substance.

The acid dissociation constant, denoted as $K_a$, quantifies the strength of an acid in an aqueous solution. It represents the equilibrium constant for the dissociation of an acid into hydronium ions ($H_3O^+$) and its conjugate base. A higher $K_a$ value indicates a stronger acid. For instance, hydrochloric acid (HCl) has a $K_a$ value of approximately $1.3 \times 10^6$, categorizing it as a strong acid.

On the other hand, the $pK_a$ of an acid is defined as the negative logarithm of its $K_a$ value:

This value serves as a measure of the acidity of the acid, where lower $pK_a$ values signify stronger acids. For example, acetic acid (CH$_3$COOH) has a $pK_a$ of approximately $4.76$, indicating that it is a weaker acid compared to HCl.

The relationship between $pK_a$ and $K_a$ is logarithmic. Specifically, a change of one unit in $pK_a$ corresponds to a tenfold variation in $K_a$. For example, an acid with a $pK_a$ of $3$ is ten times stronger than one with a $pK_a$ of $4$.

In summary, $K_a$ is a direct measure of an acid’s strength, while $pK_a$ provides a logarithmic scale to assess its acidity. The two measures are intrinsically linked, with lower $pK_a$ values indicating stronger acids.