Weak acid‑strong base titrations are a cornerstone of analytical chemistry, and the equivalence point—the moment when the amount of added base exactly neutralizes the weak acid—holds particular significance because it determines the pH jump used for indicator selection and quantitative analysis. Understanding what happens at this point, how to locate it experimentally, and why the pH is not neutral (as it would be for a strong acid‑strong base system) enables students and professionals to design accurate titrations, interpret titration curves, and apply the concept to real‑world problems such as pharmaceutical formulation and environmental monitoring.
Introduction
When a weak acid (e., acetic acid, CH₃COOH) reacts with a strong base (e.g.g.
[ \text{CH}_3\text{COOH} + \text{OH}^- \rightarrow \text{CH}_3\text{COO}^- + \text{H}_2\text{O} ]
At the equivalence point, every mole of acid initially present has been neutralized by an equivalent mole of hydroxide ions. Because of that, unlike the strong acid‑strong base case, the solution at this point contains the conjugate base of the weak acid, which hydrolyzes water to produce a mildly basic pH. Because of this, the pH at the equivalence point is >7, and the exact value depends on the acid’s dissociation constant (Kₐ) and its concentration.
Steps to Locate the Equivalence Point
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Prepare the analyte solution
- Accurately weigh a known mass of the weak acid (or pipette a known volume of a standard solution).
- Dissolve it in distilled water to a known volume, recording the initial concentration (Cₐ).
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Set up the burette
- Fill a burette with the strong base titrant (commonly NaOH) of known concentration (C_b).
- Ensure the burette is free of air bubbles and the tip is immersed in the solution.
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Add indicator or use a pH meter
- For visual titration, choose an indicator whose transition range brackets the expected equivalence‑point pH (e.g., phenolphthalein for pH ≈ 8.2–10).
- For instrumental titration, immerse a calibrated pH electrode and record pH after each increment.
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Perform the titration
- Add the base in small, consistent increments (e.g., 0.5 mL).
- Swirl the flask gently after each addition to ensure mixing.
- Note the volume of base added and the corresponding pH or color change.
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Identify the equivalence point
- Graphical method: Plot pH versus volume of base added. The equivalence point appears as the steepest slope (the inflection point).
- Derivative method: Compute ΔpH/ΔV; the maximum of this derivative marks the equivalence point.
- Indicator method: Observe the color change; the volume at which the change occurs approximates the equivalence point (less precise for weak acid‑strong base titrations).
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Calculate the acid concentration
- Use the stoichiometry: ( n_{\text{acid}} = C_b \times V_{\text{eq}} ) (where ( V_{\text{eq}} ) is the volume of base at equivalence).
- Then ( C_{\text{acid}} = \frac{n_{\text{acid}}}{V_{\text{acid solution}}} ).
Scientific Explanation
Why the pH is Not 7
In a strong acid‑strong base titration, the products are a neutral salt and water, giving pH ≈ 7 at equivalence. g.For a weak acid, the salt formed (e., sodium acetate) contains the conjugate base (acetate ion) But it adds up..
It's where a lot of people lose the thread.
[ \text{CH}_3\text{COO}^- + \text{H}_2\text{O} \rightleftharpoons \text{CH}_3\text{COOH} + \text{OH}^- ]
The equilibrium constant for this hydrolysis is the base dissociation constant ( K_b ), related to the acid’s ( K_a ) by:
[ K_b = \frac{K_w}{K_a} ]
where ( K_w = 1.0 \times 10^{-14} ) at 25 °C. The hydroxide ion concentration generated by hydrolysis determines the pH:
[ [\text{OH}^-] = \sqrt{K_b \times C_{\text{salt}}} ]
[ \text{pOH} = -\log[\text{OH}^-] \quad ; \quad \text{pH} = 14 - \text{pOH} ]
Thus, the higher the ( K_a ) (stronger weak acid), the smaller the ( K_b ) and the closer the pH is to 7. Conversely, a very weak acid (small ( K_a )) yields a larger ( K_b ) and a more basic equivalence‑point pH The details matter here..
Shape of the Titration Curve
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Initial region: pH is governed by the weak acid’s dissociation; adding base converts acid to its conjugate base, creating a buffer Worth knowing..
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Buffer region: pH changes slowly; the Henderson–Hasselbalch equation applies:
[ \text{pH} = pK_a + \log\frac{[\text{A}^-]}{[\text{HA}]} ]
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Equivalence point: All HA is converted to A⁻; the solution is no longer a buffer, and pH rises sharply due to hydrolysis.
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Post‑equivalence: Excess OH⁻ dominates; pH increases gradually with added base.
The steepness of the rise at equivalence depends on the acid’s concentration and ( K_a ); more dilute solutions or weaker acids give a less pronounced jump, making indicator selection more critical.
Choosing an Indicator
An appropriate indicator must change color within the pH range of the equivalence‑point jump. Now, for a typical weak acid (pKₐ ≈ 4–5) titrated with 0. 1 M NaOH, the equivalence‑point pH falls between 8 and 9. Phenolphthalein (pH 8.2–10.0) is therefore a common choice. For weaker acids (pKₐ > 5), the equivalence point may shift higher, necessitating indicators like thymol blue (pH 8 It's one of those things that adds up..
