1. What is pH at Half Equivalence Point?

The pH at half equivalence point in a weak acid/base titration is equal to the pKa of the acid. This relationship holds true because at this point, exactly half of the weak acid has been neutralized, creating a buffer solution where [HA] = [A⁻].

2. How Does the Calculation Work?

The calculation uses the simple formula:

Where:

• \( pH \) — Measure of acidity/alkalinity (dimensionless)
• \( pK_a \) — Acid dissociation constant (dimensionless)

Explanation: At the half-equivalence point in a titration of a weak acid with a strong base, the concentrations of the weak acid and its conjugate base are equal, making pH = pKa according to the Henderson-Hasselbalch equation.

3. Importance of pH at Half Equivalence Point

Details: This calculation is crucial in acid-base titrations as it helps determine the pKa of an unknown acid, identifies the buffering region, and is used to select appropriate indicators for titration experiments.

4. Using the Calculator

Tips: Enter the pKa value of the weak acid. The calculator will return the pH value at the half-equivalence point, which is numerically equal to the pKa.

5. Frequently Asked Questions (FAQ)

Q1: Does this relationship hold for weak bases?
A: For weak bases titrated with strong acids, at half-equivalence point, pOH = pKb, and therefore pH = 14 - pKb.

Q2: Why is pH = pKa at half-equivalence point?
A: According to the Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA]). At half-equivalence, [A⁻] = [HA], so log(1) = 0, and pH = pKa.

Q3: Is this applicable to polyprotic acids?
A: For polyprotic acids, each equivalence point has its own half-equivalence point where pH equals the pKa for that specific proton dissociation.

Q4: What are typical pKa values?
A: pKa values typically range from -10 to 50, with most common weak acids having pKa values between 2 and 12.

Q5: How accurate is this relationship?
A: This relationship is theoretically exact for ideal solutions, though activity coefficients and other non-ideal behavior may cause slight deviations in practice.