Acid–Base Titration: Apparatus Setup, pH Curves and Equivalence Points

In chemistry laboratory courses, few procedures are as common or as error-prone as titration. The setup consists of three core components:

• The Burette: A precision graduated glass tube with a stopcock, clamped vertically to a retort stand. It holds the titrant—the solution of accurately known concentration (e.g., standard 0.100 M NaOH).

• The Conical (Erlenmeyer) Flask: Placed beneath the burette tip over a white tile. It contains an accurately pipetted volume of the analyte—the solution whose concentration is unknown (e.g., vinegar or HCl)—along with a drop of chemical indicator.

• The Meniscus: Readings must be taken at eye level at the bottom of the liquid curve. Near the expected neutralization volume, the titrant must be added dropwise while constantly swirling the flask.

Plotting pH on the vertical axis against volume of added titrant on the horizontal axis produces a distinctive sigmoidal (S-shaped) curve. For a weak acid titrated with a strong base (like CH₃COOH with NaOH), the curve progresses through four critical zones:

1. Initial pH: Governed solely by the weak acid equilibrium (Ka). pH starts moderately acidic (around 2.5–3.5).

2. The Buffer Region: As OH⁻ is added, some HA is converted to conjugate base A⁻, forming an effective buffer. Exactly halfway to the equivalence point (V = ½ V_eq), the concentrations [HA] and [A⁻] are identical. By Henderson-Hasselbalch, pH = pKa + log(1) = pKa! This allows chemists to determine pKa experimentally.

3. The Equivalence Cliff: Once nearly all acid is consumed, a single drop of excess OH⁻ causes the pH to skyrocket vertically over 4–6 pH units. The center of this vertical inflection is the equivalence point.

4. Post-Equivalence Plateau: Past equivalence, excess strong base dominates, and the curve flattens asymptotically near pH 12–13.

A frequent AP Chemistry exam question tests the difference between these two terms:

• The Equivalence Point is the theoretical stoichiometric point where moles of added titrant exactly equal the moles of analyte originally present according to the balanced reaction equation.

• The Endpoint is the experimental signal—typically the moment an indicator changes color—that tells the chemist to close the burette stopcock.

A titration is accurate only if you choose an indicator whose pKa (and color-change transition interval) matches the pH of the equivalence point. For instance, phenolphthalein (turning pink at pH 8.2–10.0) is ideal for weak acid-strong base titrations whose equivalence pH is ~8.8, but methyl red would change color far too early.

Once you record the volume of titrant used (the titre), calculating analyte concentration is straightforward:

1. Calculate moles of titrant added: n_titrant = M_titrant × V_titrant(in Liters).

2. Use the stoichiometric mole ratio from the balanced equation to find moles of analyte.

3. Divide moles of analyte by the initial flask volume: M_analyte = n_analyte / V_analyte.

For polyprotic acids (like H₂SO₄ or H₃PO₄), always check the stoichiometric ratio—a diprotic acid requires twice as many moles of NaOH to reach the second equivalence point!

Reading about burettes and pH curves is nothing compared to seeing them drawn. Claryfied's chemistry engine draws laboratory glassware—burette, retort stand, flask, stirring bar—charts the pH curve point by point, and writes out the stoichiometric mole tables with synced voice explanation.

Check out our live acid-base titration lesson on Claryfied, or type in your lab data to have Claryfied calculate your unknown molarity and sketch your specific titration curve!