The comparison that answers most questions
Compare pH to pKa. When pH is below pKa, the protonated form predominates. When pH is above pKa, the deprotonated form predominates. When they are equal, the two forms are present in equal amounts.
That last case is the definition of the half-equivalence point on a titration curve, and it is the single most tested fact in this area. At half-equivalence, pH equals pKa — which is why titration curve questions so often ask you to read a pKa off a graph rather than calculate anything.
Nearly every acid–base question on the MCAT is this comparison applied to a specific molecule, plus a small amount of reasoning about what the resulting form does. Charge, solubility, membrane permeability, and enzyme activity all follow from protonation state.
Henderson–Hasselbalch without the algebra
The equation is pH = pKa + log([A⁻]/[HA]), where A⁻ is the conjugate base and HA the weak acid. On the MCAT you rarely need to solve it numerically, because the logarithm is almost always of a ratio you can handle mentally.
When the ratio is 1, the log is 0 and pH equals pKa. When the ratio is 10, the log is 1 and pH is one unit above pKa. When the ratio is 1/10, the log is −1 and pH is one unit below. When the ratio is 100, pH is two units above. Those four cases cover the overwhelming majority of questions.
Read it in the direction the question asks. Given pH and pKa, you can state the ratio of the two forms. Given the ratio, you can state the pH. Given the pH and the desired ratio, you can select the right buffer from a table of pKa values — which is the standard laboratory-design question.
A useful consequence: a one-unit gap between pH and pKa already means a ten-to-one imbalance, so "predominates" understates it. At physiological pH of 7.4, a group with a pKa of 4.4 is deprotonated by a ratio of a thousand to one, which is why such groups are treated as fully charged.
What makes a buffer work
A buffer is a solution containing appreciable amounts of both a weak acid and its conjugate base. Added acid is consumed by the conjugate base; added base is consumed by the weak acid. Because both reservoirs exist, pH changes far less than it would in unbuffered solution.
Buffering capacity is greatest when pH is close to pKa, and it is maximal at pH equal to pKa, where the two forms are equimolar and neither reservoir is close to exhausted. The working range is conventionally pKa plus or minus one pH unit — outside that, one form is outnumbered ten to one and the buffer is nearly spent in that direction.
This produces the standard selection question: to buffer at a target pH, choose the weak acid whose pKa is nearest that pH. A buffer with a pKa three units away is not a weak buffer; it is effectively not a buffer at all.
Strong acids and bases cannot buffer. A strong acid dissociates completely, so no reservoir of undissociated acid exists to absorb added base. This is also why strong-acid titration curves have no flat region before the equivalence point, while weak-acid curves do — that flat region is the buffering zone, centered on the pKa.
Reading a titration curve
| Point on the curve | What is true there | Commonly tested |
|---|---|---|
| Start | Essentially all HA | Initial pH depends on Ka |
| Half-equivalence | [HA] = [A⁻], so pH = pKa | Read pKa directly off the graph; maximal buffering |
| Buffering region | Both forms present; curve is flat | Roughly pKa ± 1 |
| Equivalence point | Moles of base added equal moles of acid initially present | For a weak acid titrated with strong base, pH here is above 7 |
| Beyond equivalence | Excess titrant dominates | pH approaches that of the titrant |
The equivalence point is where the most errors occur, because students expect pH 7. That holds only for a strong acid titrated with a strong base. Titrating a weak acid with a strong base leaves the conjugate base A⁻ in solution at equivalence, and A⁻ is itself basic, so the pH is above 7. The mirror case — a weak base titrated with strong acid — gives an equivalence pH below 7.
A polyprotic acid produces one buffering region and one equivalence point per ionizable proton, so a diprotic acid gives a curve with two plateaus and two steep jumps. Each plateau's midpoint gives one pKa.
The bicarbonate system and physiology questions
The blood's principal buffer is carbonic acid and bicarbonate, linked by the reaction CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻. Its importance is not its pKa, which at roughly 6.1 is not especially close to blood pH of 7.4, but that it is an open system: both components are independently regulated.
The lungs control CO₂ by ventilation, and the kidneys control bicarbonate by reabsorption and excretion. Because the acid component can be exhaled, the system buffers far more effectively than its pKa alone would predict.
Apply Le Châtelier's principle to answer most questions here. Hypoventilation retains CO₂, pushing the equilibrium right, raising H⁺, and lowering pH — respiratory acidosis. Hyperventilation blows off CO₂, pulling the equilibrium left and raising pH — respiratory alkalosis. A metabolic acidosis, such as accumulating lactic acid, consumes bicarbonate and is compensated by increased ventilation to lower CO₂.
Note the direction of compensation: a respiratory problem is compensated renally over hours to days, and a metabolic problem is compensated respiratorily within minutes. Questions frequently turn on which system caused the disturbance and which is compensating.
The common mistakes
Confusing pKa with pH. pKa is a fixed property of a molecule; pH is a property of the solution. The question is always how they compare, and neither one alone tells you the protonation state.
Reversing the direction of the comparison. Low pH means high H⁺ concentration, which drives protonation. It is worth deriving this once from Le Châtelier rather than memorizing it, because the memorized version reverses under pressure.
Assuming the equivalence point is at pH 7. True only for strong acid with strong base. Weak acid with strong base gives an equivalence pH above 7.
Treating a lower pKa as a weaker acid. Lower pKa means stronger acid — more willing to donate a proton, and a more stable conjugate base. The inverse relationship trips people up because the numbers run opposite to the intuition.
Forgetting that adding a buffer's conjugate base changes the ratio, not the pKa. Buffer-preparation questions ask you to manipulate the ratio to land on a target pH; the pKa is a constant you are working around.
Frequently asked questions
When does pH equal pKa?
When the concentrations of the weak acid and its conjugate base are equal. On a titration curve this is the half-equivalence point, which is also where buffering capacity is greatest.
How do I use Henderson–Hasselbalch on the MCAT?
Usually without calculating. A base-to-acid ratio of 1 puts pH at pKa; a ratio of 10 puts it one unit above; a ratio of 1/10 puts it one unit below. Those cases cover most questions.
How do I choose the right buffer for a target pH?
Select the weak acid whose pKa is closest to the target pH. Effective buffering runs roughly one pH unit either side of the pKa, so a buffer chosen further away provides very little capacity.
Why is the equivalence point of a weak acid titration above pH 7?
Because at equivalence all the weak acid has been converted to its conjugate base, which is itself basic and accepts protons from water. The resulting solution is basic.
Why is the bicarbonate buffer effective at pH 7.4 despite a pKa near 6.1?
Because it is an open system. CO₂ is removed by the lungs and bicarbonate is regulated by the kidneys, so both components can be adjusted independently rather than being fixed in a closed solution.
Related Verbloom guides
Want CARS reading to feel methodical?
Verbloom drills the argument-first reading that CARS rewards — main idea, author attitude, structure, and supported inference — in short, focused sessions.