Titration of a weak acid and strong base is a fundamental laboratory technique used to determine the concentration of an acidic solution by reacting it with a standardized strong base solution. This method provides valuable insight into acid‑base equilibria, buffer behavior, and the calculation of dissociation constants, making it a cornerstone of analytical chemistry education and quality‑control laboratories.
Theory Behind the Titration
When a weak acid (HA) is titrated with a strong base (such as NaOH), the reaction proceeds according to the simple stoichiometry
[ \text{HA} + \text{OH}^- \rightarrow \text{A}^- + \text{H}_2\text{O} ]
Unlike a strong‑acid/strong‑base titration, the pH does not jump abruptly at the equivalence point because the conjugate base (A⁻) formed in the solution hydrolyzes water, producing a basic solution. Understanding the shape of the pH curve requires familiarity with several key concepts Simple, but easy to overlook..
Acid Dissociation Constant (Ka)
The intrinsic strength of the weak acid is expressed by its acid dissociation constant
[ K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]} ]
or, more conveniently, its negative logarithm
[ \text{p}K_a = -\log K_a ]
Knowing the pKa allows prediction of the pH at various points during the titration, especially at the half‑equivalence point where ([HA] = [A^-]) and pH = pKa And it works..
Buffer Region and Henderson‑Hasselbalch Equation
Before the equivalence point, the solution contains appreciable amounts of both the weak acid and its conjugate base, forming a buffer. The pH of this buffer can be estimated with the Henderson‑Hasselbalch equation
[ \text{pH} = \text{p}K_a + \log\frac{[\text{A}^-]}{[\text{HA}]} ]
As base is added, the ratio ([\text{A}^-]/[\text{HA}]) increases, causing a gradual rise in pH. This region appears as a relatively flat slope on the titration curve Easy to understand, harder to ignore..
Equivalence Point and Beyond
At the equivalence point, moles of added OH⁻ equal the initial moles of HA. All HA has been converted to A⁻, and the solution contains only the conjugate base (plus the spectator cation from the strong base). The pH is therefore governed by the hydrolysis reaction
[ \text{A}^- + \text{H}_2\text{O} \rightleftharpoons \text{HA} + \text{OH}^- ]
with a base hydrolysis constant
[ K_b = \frac{K_w}{K_a} ]
Because of this, the pH at the equivalence point is >7 and can be calculated from
[ [\text{OH}^-] = \sqrt{K_b \cdot C_{A^-}} ]
where (C_{A^-}) is the concentration of the conjugate base after dilution. After the equivalence point, excess strong base dominates, and the pH rises sharply, reflecting the direct contribution of ([\text{OH}^-]) from the added NaOH.
Experimental Procedure
Materials and Reagents
- Weak acid solution (e.g., acetic acid, unknown concentration)
- Standardized strong base (typically 0.100 M NaOH)
- pH meter with calibrated electrode or a suitable indicator (phenolphthalein, bromothymol blue)
- Burette (50 mL) and stand
- Erlenmeyer flask (100–250 mL)
- Distilled water for rinsing
- Magnetic stirrer and stir bar (optional)
Step‑by‑Step Procedure
- Prepare the burette – Rinse with distilled water, then with a small amount of the NaOH solution. Fill the burette with the standardized NaOH, noting the initial volume to the nearest 0.01 mL.
- Transfer the acid – Pipette a known volume (e.g., 25.00 mL) of the weak acid solution into a clean Erlenmeyer flask. Add ~20 mL of distilled water to ensure adequate mixing.
- Add indicator – If using an indicator, add 2–3 drops of phenolphthalein (colorless in acid, pink in base). For pH‑meter monitoring, skip this step.
- Begin titration – Place the flask under the burette tip, start stirring, and add NaOH in small increments (0.5 mL). After each addition, record the burette reading and the corresponding pH (or observe color change).
- Approach the equivalence point – As the pH begins to rise more rapidly, decrease the increment size to 0.1 mL or less to pinpoint the steep portion of the curve.
- Pass the equivalence point – Continue adding NaOH until a stable, high pH (≈11–12) is reached or the indicator shows a permanent color change. Record the final burette volume.
- Repeat – Perform at least two additional titrations to ensure reproducibility; average the volumes of NaOH used at the equivalence point.
Data Collection and Analysis
Plotting the pH Curve
Using the recorded volumes of NaOH (x‑axis) and measured pH values (y‑
-axis) to visualize the titration profile. The equivalence point is identified as the point where the pH changes most abruptly, corresponding to the stoichiometric neutralization of the weak acid by the strong base. The volume of NaOH at this inflection point is recorded as ( V_{\text{eq}} ).
Calculation of Weak Acid Concentration
At the equivalence point, the moles of NaOH added are stoichiometrically equal to the moles of weak acid initially present:
[ n_{\text{HA}} = n_{\text{NaOH}} = M_{\text{NaOH}} \cdot V_{\text{eq}} ]
where ( M_{\text{NaOH}} ) is the molarity of the standardized base. The concentration of the weak acid is then calculated using:
[ C_{\text{HA}} = \frac{n_{\text{HA}}}{V_{\text{acid}}} ]
with ( V_{\text{acid}} ) being the initial volume of the weak acid solution pipetted into the flask.
Theoretical pH at Equivalence Point
The experimental pH at equivalence can be compared to the theoretical value derived from the hydrolysis reaction of the conjugate base (( \text{A}^- )). Using the earlier expression for ( [\text{OH}^-] ):
[ [\text{OH}^-] = \sqrt{K_b \cdot C_{A^-}} ]
where ( C_{A^-} ) is the concentration of the conjugate base after dilution (calculated as ( C_{A^-} = \frac{n_{\text{HA}}}{V_{\text{total}}} ), with ( V_{\text{total}} = V_{\text{acid}} + V_{\text{eq}} )). The theoretical pH is then:
[ \text{pH} = 14 + \log{[\text{OH}^-]} ]
This comparison validates the assumption that the conjugate base hydrolyzes to produce a basic solution.
Buffer Regions and Titration Curves
Before equivalence, the solution acts as a buffer, with the pH governed by the Henderson-Hasselbalch equation:
[ \text{pH} = \text{p}K_a + \log{\left( \frac{[\text{A}^-]}{[\text{HA}]} \right)} ]
Plotting ( \log{\left( \frac{[\text{A}^-]}{[\text{HA}]} \right)} ) against the volume of NaOH added allows for the determination of ( \text{p}K_a ) from the slope of the buffer region. After equivalence, the excess ( \text{OH}^- ) from NaOH dominates, and the pH is calculated directly from:
[ [\text{OH}^-] = \frac{M_{\
