An acid-base indicator is a weak acid or weak base that signals the pH at the endpoint of a titration. The endpoint occurs just after the equivalence point, marked by a color change in the indicator. This change signifies that the weak form of the indicator has transformed into its conjugate form, which exhibits a different color. For instance, if the indicator is a weak acid, it will change color as it converts to its conjugate base form after the endpoint is reached. Conversely, if a weak base is used as the indicator, it gains a proton (H⁺) and shifts to its weak acid form, resulting in a distinct color.

The effectiveness of an indicator in a titration is determined by how closely its pK_a value aligns with the pH at the equivalence point. The pH range of an indicator is typically expressed as pH = pK_a ± 1, which is similar to the buffer range concept. In a titration curve, the color of the indicator changes as the pH shifts. Initially, a solution may appear purple, indicating a higher concentration of the weak acid. As a base is added, the solution transitions to a pinkish hue at the endpoint, where the concentrations of the weak acid and its conjugate base are equal. Beyond the equivalence point, the solution may turn reddish due to excess strong base, indicating a higher concentration of the conjugate base.

In terms of pH and pK_a relationships, the following can be observed: in the purple region, pH = pK_a - 1; at the endpoint, pH = pK_a; and beyond the endpoint, pH = pK_a + 1. Common acid-base indicators include methyl orange, which operates best in the pH range of 3.3 to 4.5, transitioning from red to yellow. Another example is thymol blue, which has two pH ranges: 1.2 to 2.8 (red to yellow) and 8 to 9.2 (yellow to blue). These indicators demonstrate the various color transitions that occur as they shift between their original and conjugate forms, highlighting their utility in titrations.