In a standard acid-base titration curve, the x-axis displays the volume of the titrant added from a buret, while the y-axis shows the pH of the analyte solution in the flask. As the titrant is gradually mixed into the solution, the pH changes. Initially, the pH changes slowly as the chemical species in the flask neutralize the incoming drops.
However, near the equivalence point, which is the exact moment where the moles of added titrant perfectly equal the moles of the initial analyte, the graph features a very steep, nearly vertical slope. This dramatic jump in pH is the hallmark signature of a titration curve.
The exact shape and features of a titration curve depend heavily on the strengths of the acids and bases involved. When a strong acid is titrated with a strong base, the curve starts at a very low pH, stays relatively flat, and then spikes drastically through a neutral equivalence point of exactly pH 7.
In contrast, the titration of a weak acid with a strong base produces a curve with a distinctive buffering region early on. In this region, the pH rises slowly because the weak acid and its newly formed conjugate base act together as a buffer. Furthermore, the equivalence point for a weak acid and strong base is always located at a pH greater than 7 due to the basic nature of the conjugate base that remains.
Titration curves are exceptionally valuable tools in analytical chemistry for multiple practical reasons. First, they allow chemists to precisely determine the unknown concentration of a sample by locating the exact midpoint of the steep vertical section. Second, for weak acids or weak bases, the curve can be used to find the acid dissociation constant (Ka).
At the exact halfway mark to the equivalence point, the pH of the solution is mathematically equal to the pKa of the acid. Finally, the curve helps in selecting the proper chemical indicator, as the ideal indicator must change color entirely within the steep pH range shown on the graph.