How do Acid-Base Titration and the Henderson-Hasselbalch Equation work?
Acid-base titration is a foundational quantitative chemical analysis technique used to determine the concentration of an acid analyte ($C_a$) by gradually adding a standardized base titrant ($C_b$, such as $0.10\text{ M NaOH}$) from a burette. The equivalence point is reached when the moles of added $\text{OH}^-$ exactly equal the initial moles of acid: $C_a \cdot V_a = C_b \cdot V_{\text{eq}}$.
When titrating a strong acid ($\text{HCl}$) with a strong base ($\text{NaOH}$), the resulting salt ($\text{NaCl}$) is neutral and the equivalence point sits at exactly $\text{pH} = 7.00$ at $25\text{ °C}$. By contrast, when titrating a weak acid such as acetic acid ($\text{CH}_3\text{COOH}$, $\text{p}K_a = 4.76$), prior to equivalence a conjugate acid-base buffer forms, governed by the Henderson-Hasselbalch Equation: $\text{pH} = \text{p}K_a + \log_{10}\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)$.
At the half-equivalence point ($V_b = \frac{1}{2}V_{\text{eq}}$), exactly half the acid is neutralized ($[\text{A}^-] = [\text{HA}]$), so $\text{pH} = \text{p}K_a = 4.76$. At the equivalence point ($V_b = V_{\text{eq}}$), hydrolysis of the acetate ion raises the pH to $\sim 8.72$.
Step-by-step guided experiment
1. The steep equivalence jump: With $0.10\text{ M HCl}$ ($V_{\text{eq}} = 25.0\text{ mL}$), slide $V_b$ to $24.8\text{ mL}$ ($\text{pH} = 3.40$, colorless with phenolphthalein). Add just $0.4\text{ mL}$ more to reach $25.2\text{ mL}$: the pH leaps over 7 units to $10.60$ and the flask turns bright magenta pink.
2. Weak acid buffer region: Switch the analyte to Acetic Acid ($\text{CH}_3\text{COOH}$) and set $V_b = 12.5\text{ mL}$ (half of $25.0\text{ mL}$). Verify on the digital readout that $\text{pH} = 4.76 = \text{p}K_a$.
Frequently Asked Questions (FAQ)
Why is the equivalence point pH greater than 7 when titrating a weak acid with a strong base?
When a weak acid such as acetic acid ($\text{CH}_3\text{COOH}$) is neutralized by $\text{NaOH}$, at the equivalence point all the acid has been converted into its conjugate base ($\text{CH}_3\text{COO}^-$), which partially hydrolyzes in water to produce $\text{OH}^-$ ions, raising the pH to $\sim 8.72$.
What happens at the half-equivalence point of a weak acid?
When exactly half the volume of base required for equivalence has been added, the concentrations of weak acid $[\text{HA}]$ and conjugate base $[\text{A}^-]$ are equal. By the Henderson-Hasselbalch equation, at that exact point $\text{pH} = \text{p}K_a$.
Which chemical indicator should be chosen for a titration?
Choose an indicator whose color transition range matches the steep pH jump at the equivalence point: phenolphthalein (transitions at pH $8.2 – 10.0$) is ideal for both weak and strong acids titrated with $\text{NaOH}$.
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