Problem: A buffer contains 0.20 M acetic acid (pKa = 4.756) and 0.30 M sodium acetate. Find the pH using the Henderson-Hasselbalch equation.

Problem: You need an acetate buffer at pH 5.0 (pKa = 4.756) with total concentration 0.5 M. Find [CH₃COOH] and [CH₃COO⁻] using the Henderson-Hasselbalch equation.

Problem: A buffer at pH 7.2 contains 0.40 M H₂PO₄⁻ and 0.60 M HPO₄²⁻. Find pKa using the Henderson-Hasselbalch equation.

Problem: An ammonia buffer (pKa of NH₄⁺ = 9.25) contains 0.15 M NH₃ and 0.05 M NH₄Cl. Find the pH using the Henderson-Hasselbalch equation.

Note: even though NH₃ is a weak base, the Henderson-Hasselbalch equation still applies using the pKa of its conjugate acid, NH₄⁺, with NH₃ playing the role of [A⁻] (base) and NH₄⁺ playing the role of [HA] (acid).

Problem: Calculate the pH of a buffer with pKa = 4.756, [A⁻] = 0.001 M, [HA] = 1.0 M using the Henderson-Hasselbalch equation.

The ratio [A⁻]/[HA] = 0.001 is far outside the effective buffering range of 0.1 to 10 (i.e., pH within ±1 of pKa). This is not an effective buffer — the Henderson-Hasselbalch equation still produces a numeric answer, but the solution will have almost no actual buffering capacity against added acid or base.

Problem: 500 mL of 0.1 M acetate buffer is at pH 4.756. How many mL of 1 M NaOH must be added to shift it to pH 5.0? (pKa = 4.756)

Problem: A Tris buffer is made at 25°C to pH 8.0 (pKa = 8.072 at 25°C). What is the pH at 37°C, given pKa changes by −0.031 per °C?

This is why Tris buffers are notorious in biochemistry labs — a Henderson-Hasselbalch calculation done at room temperature can be off by nearly half a pH unit once the buffer is used at body temperature, because Tris’s pKa is unusually temperature-sensitive compared to phosphate or acetate buffers.

Problem: Calculate the maximum buffer capacity of a 0.2 M phosphate buffer.

This means the buffer can neutralize about 0.115 mol of strong acid per liter before the pH shifts by a full unit — but only when the buffer is at pH = pKa. Away from pKa, buffer capacity drops off, which connects directly to the effective range discussed in Problem 5.

Problem: Design a phosphate buffer at pH 7.4 with total [phosphate] = 50 mM in 1 L, using NaH₂PO₄ (MW = 120.0 g/mol) and Na₂HPO₄ (MW = 142.0 g/mol). pKa = 7.198.

This is a real laboratory buffer preparation calculation — the Henderson-Hasselbalch equation gives the molar ratio, and multiplying by molecular weight converts it into a practical recipe you can actually weigh out on a balance.

Problem: If [A⁻] in a buffer is increased while [HA] is held constant, does the pH increase, decrease, or stay the same?

This type of reasoning-only Henderson-Hasselbalch practice problem is common on the MCAT — you’re expected to predict the direction of a pH shift without plugging in any numbers at all, just from understanding the structure of the equation.

Problem: A citrate buffer has pKa₂ = 4.761. At pH 4.5, [H₂Cit⁻] = 0.08 M. Find [HCit²⁻] using the Henderson-Hasselbalch equation.

Citrate is a triprotic acid with three separate pKa values — always confirm you’re using the correct pKa (here, pKa₂) that corresponds to the specific conjugate acid/base pair given in the problem.

Problem: 20 mL of 0.1 M acetic acid (pKa = 4.756) is titrated with 0.1 M NaOH. Find the pH after adding 8 mL of NaOH.

Key insight: notice that the total volume (28 mL) never actually appears in the calculation, because it cancels out identically in both the numerator and denominator of the ratio. This is a detail many students miss — in titration problems, you can use moles directly instead of converting to concentrations, since the Henderson-Hasselbalch equation only needs the ratio, not the absolute concentrations.