A buffer is prepared with 0.10 M acetic acid and 0.10 M sodium acetate (pKa = 4.76). A small amount of HCl is added. What happens to the pH?
ApH stays exactly at 4.76 — the buffer is at its sweet spot and completely neutralizes the acid
BpH drops below 3 immediately, because HCl is a strong acid
CpH drops slightly as acetate ions react with added H⁺, converting some conjugate base to weak acid
DpH increases because the buffer absorbs the added acid
Buffers resist pH change but do not prevent it entirely. The added H⁺ reacts with acetate (A⁻) via Le Chatelier shift, consuming some conjugate base and producing more acetic acid — so pH drops slightly. Option A is the most common misconception: pH = pKa only when [A⁻]/[HA] = 1, but adding acid shifts that ratio, causing a small but real pH decrease.
Question 2 Multiple Choice
A researcher needs a buffer with twice the capacity of their current 0.10 M acetate buffer at pH 4.76. They prepare a new buffer at the same pH ratio but with 0.20 M concentrations of both components. How does the new buffer compare?
ASame capacity — pH and ratio are identical, so resistance to change is the same
BGreater capacity — it can absorb twice as much added acid or base before the buffering fails
CLesser capacity — higher concentrations shift the equilibrium away from the optimal ratio
DGreater capacity, but only for added acid, not added base
Buffer capacity depends on the absolute concentrations of weak acid and conjugate base, not just their ratio. Both buffers have the same pH (ratio determines pH via Henderson-Hasselbalch), but the 0.20 M buffer has twice the moles of each component available to absorb added acid or base before being overwhelmed.
Question 3 True / False
A buffer prepared with a weak acid and its conjugate base at a 10:1 ratio (more weak acid than conjugate base) will have a pH below the pKa of the weak acid.
TTrue
FFalse
Answer: True
pH = pKa + log([A⁻]/[HA]). With [A⁻]/[HA] = 1/10, log(1/10) = −1, so pH = pKa − 1. More weak acid than conjugate base pushes pH below pKa. This also means the buffer has less capacity to absorb added acid (conjugate base is nearly depleted in that direction).
Question 4 True / False
Two buffer solutions with the same weak acid/conjugate base ratio will resist pH change equally well, regardless of their absolute concentrations.
TTrue
FFalse
Answer: False
The ratio determines pH via Henderson-Hasselbalch, but the absolute concentrations determine buffer capacity — how much acid or base can be absorbed before the buffer fails. A 1.0 M acetate buffer and a 0.01 M acetate buffer at the same ratio have identical pH but the former can absorb 100 times more added acid or base before being exhausted.
Question 5 Short Answer
Why must an effective buffer contain significant amounts of both the weak acid and its conjugate base, rather than having one component greatly exceed the other?
Think about your answer, then reveal below.
Model answer: With both components present in comparable amounts, the buffer has capacity in both directions: conjugate base absorbs added H⁺ (Le Chatelier shifts equilibrium toward HA), and weak acid donates H⁺ to absorb added OH⁻ (shifts toward A⁻). If one component is nearly depleted, the system loses the ability to resist changes in that direction. Maximum buffering capacity and pH stability occur at the 1:1 ratio (pH = pKa) because the system has equal reserves in both directions.
This is the mechanistic heart of buffering. The Henderson-Hasselbalch equation quantifies it: the buffer works best within ±1 pH unit of pKa, where neither component is below ~10% of the other. Outside that range, one reservoir is so small that adding even a little acid or base overwhelms it, and pH changes rapidly.