Acid Base Titration Weak Acid Strong Base
An acid‑base titration involving a weak acid and a strong base is a fundamental laboratory technique used to determine the concentration of an unknown acidic solution, to study acid dissociation constants (Ka), and to understand buffer behavior. In this type of titration, the titrant (usually NaOH or KOH) is a strong base that dissociates completely, while the analyte is a weak acid such as acetic acid, formic acid, or phosphoric acid’s first proton. Because the weak acid only partially ionizes, the pH changes gradually at the start of the titration, exhibits a noticeable buffer region, and shows a relatively sharp rise near the equivalence point. Understanding the shape of the titration curve, selecting an appropriate indicator, and performing accurate calculations are essential for obtaining reliable results Worth keeping that in mind..
Theoretical Background
When a weak acid (HA) reacts with a strong base (OH⁻), the neutralization reaction proceeds as follows:
[ \text{HA} + \text{OH}^- \rightarrow \text{A}^- + \text{H}_2\text{O} ]
The strong base supplies hydroxide ions that quantitatively remove protons from the weak acid, forming its conjugate base (A⁻). Since HA does not fully dissociate in solution, the initial pH is higher than that of a strong acid at the same concentration. As base is added, the solution contains a mixture of HA and its conjugate base A⁻, creating a buffer system described by the Henderson–Hasselbalch equation:
[ \text{pH} = \text{p}K_a + \log\frac{[\text{A}^-]}{[\text{HA}]} ]
At the equivalence point, moles of OH⁻ added equal the initial moles of HA. All HA has been converted to A⁻, and the solution now contains only the conjugate base, which hydrolyzes water to produce OH⁻:
[ \text{A}^- + \text{H}_2\text{O} \rightleftharpoons \text{HA} + \text{ + OH}^-} ]
So naturally, the pH at the equivalence point is basic (typically pH 8–10 for monoprotic weak acids). The exact pH can be calculated from the Kb of A⁻ (Kb = Kw/Ka) and the concentration of the salt formed Surprisingly effective..
Experimental Procedure
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Prepare Solutions
- Standardize the strong base (e.g., 0.100 M NaOH) using a primary standard such as potassium hydrogen phthalate (KHP).
- Prepare the weak acid solution of unknown concentration (e.g., acetic acid) in a clean volumetric flask.
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Set Up the Titration Apparatus
- Rinse a burette with the standardized base, then fill it, recording the initial volume.
- Pipette a known volume (usually 25.00 mL) of the weak acid into an Erlenmeyer flask.
- Add 2–3 drops of a suitable indicator (see section below).
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Perform the Titration
- Place the flask on a white tile under the burette tip.
- Add base rapidly at first, then slow down as the color change approaches.
- Near the endpoint, add base dropwise, swirling the flask after each addition.
- Record the burette reading when a persistent color change (usually pink for phenolphthalein) appears for at least 30 seconds.
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Repeat
- Conduct at least three titrations to obtain concordant results (within 0.05 mL).
Calculations
Determining Moles of Base Used
[ n_{\text{OH}^-} = C_{\text{base}} \times V_{\text{base}} ;(\text{L}) ]
where (C_{\text{base}}) is the molarity of the standardized base and (V_{\text{base}}) is the volume delivered at the endpoint.
Moles of Weak Acid
Because the reaction is 1:1,
[ n_{\text{HA}} = n_{\text{OH}^-} ]
Concentration of the Weak Acid
[ C_{\text{HA}} = \frac{n_{\text{HA}}}{V_{\text{HA}} ;(\text{L})} ]
with (V_{\text{HA}}) the pipetted volume of the acid solution.
Finding pKa from the Half‑Equivalence Point
At half‑equivalence, ([\text{HA}] = [\text{A}^-]), so the Henderson–Hasselbalch equation simplifies to:
[ \text{pH}_{\frac{1}{2}} = \text{p}K_a ]
Thus, the pH recorded when exactly half the volume of base needed for equivalence has been added gives a direct estimate of the acid’s pKa That alone is useful..
pH at Equivalence Point (Optional)
If the concentration of the salt (A⁻) at equivalence is known, the pH can be calculated:
[ [\text{A}^-] = \frac{n_{\text{HA}}}{V_{\text{total}}} ] [ K_b = \frac{K_w}{K_a} ] [ [\text{OH}^-] = \sqrt{K_b \times [\text{A}^-]} ] [ \text{pOH} = -\log[\text{OH}^-], \quad \text{pH} = 14 - \text{pOH} ]
Choice of Indicator
Because the equivalence point lies in the basic range, an indicator that changes color around pH 8–10 is required. Common choices include:
- Phenolphthalein (pH 8.2–10.0) – colorless to pink; widely used for weak acid–strong base titrations.
- Thymol blue (second transition, pH 8.0–9.6) – yellow to blue.
- Alizarin yellow R (pH 10.0–12.0) – yellow to red; useful for very weak acids.
Using an indicator with a transition range that does not encompass the equivalence point leads to systematic error (either premature or delayed endpoint detection) It's one of those things that adds up..
Sources of Error and How to Minimize Them
| Error Type | Origin | Mitigation |
|---|---|---|
| Incomplete mixing | Local excess of base causing overshoot | Swirl flask continuously; add base dropwise near endpoint |
| Carbonate contamination | Atmospheric CO₂ dissolves in base, forming carbonate that consumes acid | Prepare fresh NaOH solution; store in a tightly sealed bottle; optionally bubble nitrogen through solution |
| Indicator impurity | Indicator itself may be acidic/basic | Use high‑purity indicator; add only a few drops |
| Temperature variation | Kw and Ka are temperature‑dependent | Perform titration at a constant temperature (≈25 °C) or apply temperature corrections |
| Burette reading parallax | Misreading the meniscus | Read at eye level; use a burette with a white backing card |
| Endpoint overshoot | Adding too much base after color change | Approach endpoint slowly; confirm persistence of color for ≥30 s |
Frequently Asked Questions
Q: Why does the pH at the equivalence point exceed 7 for a weak acid–strong base titration?
