Acids and Bases - pH curves (HL only)

The equivalence point in a titration is the point where the amount of titrant added is stoichiometrically equal to the amount of analyte present. The $pH$ at this point depends on the nature of the salts formed.

Strong Acid (SA) - Strong Base (SB) titration: The $pH$ at the equivalence point is exactly $7.0$ (at $298K$) because the resulting salt does not undergo hydrolysis.

Weak Acid (WA) - Strong Base (SB) titration: The $pH$ at the equivalence point is $> 7$ because the conjugate base ($A^-$) reacts with water (hydrolysis): $A^-(aq) + H_2O(l) ightleftharpoons HA(aq) + OH^-(aq)$.

Strong Acid (SA) - Weak Base (WB) titration: The $pH$ at the equivalence point is $< 7$ because the conjugate acid ($BH^+$) reacts with water: $BH^+(aq) + H_2O(l) ightleftharpoons B(aq) + H_3O^+(aq)$.

The half-equivalence point is reached when half the volume of titrant required for equivalence has been added. At this point, for a weak acid titration, $[HA] = [A^-]$, and therefore $pH = pK_a$.

The buffer region is the range of the $pH$ curve (usually for WA-SB or SA-WB) where the $pH$ changes very slowly despite the addition of titrant. This occurs when both the weak species and its salt are present in significant concentrations.

Acid-Base indicators are weak acids or bases that change color over a specific $pH$ range. For an indicator to be effective, its $pK_{in}$ should be as close as possible to the $pH$ at the equivalence point, or fall within the steep vertical section of the curve.