Acids differ in how completely they give up a proton in water. That tendency is measured by the acid ionization constant, Ka, which is large for strong acids and small for weak ones. Because Ka values span many orders of magnitude, chemists convert them to a simpler scale by taking pKa equal to the negative logarithm of Ka. This turns awkward numbers like 1.8 times 10 to the minus 5 into an easy value of about 4.74.
Because of the negative sign in the definition, the relationship between pKa and strength is inverted. A small Ka means a large pKa and a weak acid, while a large Ka means a small or even negative pKa and a strong acid. Acetic acid has a pKa near 4.76 and is weak, while hydrochloric acid has a pKa below zero and is strong. A difference of one pKa unit represents a tenfold difference in the ionization constant.
The pKa has a direct physical meaning tied to pH. When the pH of a solution equals the pKa of an acid, exactly half of the acid is in its protonated form and half is in its deprotonated form. This is why the pKa marks the center of a buffer‘s most effective range and the midpoint of a titration curve.
The idea extends beyond acids. A base can be described by its pKb, and a conjugate acid-base pair follows the simple relationship pKa plus pKb equals 14 in water at 25 degrees Celsius. Chemists often list the pKa of the conjugate acid to compare base strength on a single scale.
Knowing pKa values helps predict the direction of proton-transfer reactions, since a proton tends to move from the stronger acid to form the weaker one.