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IB Chemistry Buffer Solutions (HL): Questions Answered
Answered by RevisionPrep's IB Educators
Buffer solutions trip up more HL Chemistry students than any other part of Reactivity 3 — not because the chemistry's hard, but because the calculations punish sloppy log work. Here's how to actually answer these questions, calculate buffer pH correctly, and avoid the mistakes examiners see every session. Answered by RevisionPrep's IB Educators.
Understanding Buffer Solutions
What is a buffer solution in IB Chemistry?
A buffer solution resists a change in pH when you add small amounts of acid or base. In IB Chemistry it's a mixture of a weak acid and its conjugate base (or a weak base and its conjugate acid) in comparable, significant concentrations — usually made by partially neutralising a weak acid or mixing it with a salt of its conjugate base.
According to the IB Chemistry guide (first assessed 2025), buffers sit inside Reactivity 3.4, the HL-only pH curves sub-topic — SL students never see this content. Two classic examples I use in class: ethanoic acid/sodium ethanoate (acidic buffer, pH below 7) and ammonia/ammonium chloride (basic buffer, pH above 7).
How do you answer buffer solutions questions in IB Chemistry?
Work out first whether the question wants an explanation or a calculation. For 'explain' marks, write the equilibrium equation and describe which species mops up added H+ or OH-. For 'calculate' marks, set up an ICE table, then substitute the acid and conjugate base amounts into the Henderson–Hasselbalch equation.
- Write the relevant equilibrium (e.g. CH3COOH ⇌ CH3COO- + H+). 2. State which component neutralises added acid, which neutralises added base. 3. Find moles of acid and conjugate base from the question. 4. Substitute into pH = pKa + log([A-]/[HA]). Quick tip: because both species share the same volume, you can use the mole ratio directly — you don't need separate concentrations, which saves time under exam pressure.
How do buffer solutions work to resist pH change?
A buffer contains a reserve of weak acid to neutralise added OH- and a reserve of conjugate base to neutralise added H+. Because both species are present in large, similar amounts, the equilibrium shifts slightly to absorb the addition rather than letting free H+ or OH- concentration jump, so pH barely moves.
This is Le Chatelier's principle in action on the buffer equilibrium. But capacity is finite — add too much acid and you exhaust the conjugate base reserve. Buffers generally work well within about one pH unit either side of the pKa (or pKb for basic buffers); outside that range, pH starts moving fast.
What's the difference between an acidic buffer and a basic buffer?
An acidic buffer combines a weak acid with a salt of its conjugate base and holds a pH below 7 — ethanoic acid/sodium ethanoate is the standard IB example. A basic buffer combines a weak base with a salt of its conjugate acid and holds a pH above 7, like ammonia/ammonium chloride.
The maths mirrors each other: acidic buffers use pH = pKa + log([A-]/[HA]); basic buffers use pOH = pKb + log([BH+]/[B]), then pH = 14 − pOH at 25°C. Mixing these two equations up under exam pressure is a very common slip.
Buffer Calculations & Worked Examples
How do you calculate the pH of a buffer solution?
Use the Henderson–Hasselbalch equation: pH = pKa + log([conjugate base]/[weak acid]). Find pKa from -log(Ka), then substitute the equilibrium concentrations (or moles, since volume cancels) of the acid and its conjugate base. This is the standard method examiners expect for any 'calculate the pH of a buffer' question on Paper 2.
Worked example: a buffer contains 0.20 mol CH3COOH and 0.30 mol CH3COONa in 1 dm3 of solution. Ka(CH3COOH) = 1.8 × 10^-5, so pKa = 4.74. pH = 4.74 + log(0.30/0.20) = 4.74 + 0.18 = 4.92. Common mistake: plugging in Ka instead of pKa — always take the -log first.
How do you prepare a buffer solution with a specific pH?
Pick a weak acid whose pKa is close to your target pH (within about one unit), then work backwards through Henderson–Hasselbalch to find the ratio of conjugate base to acid you need. Mix those amounts — either two salts, or partially neutralise the acid with strong base — in one solution.
Worked example: target pH 5.00 using ethanoic acid, pKa 4.74. 5.00 = 4.74 + log([A-]/[HA]), so log([A-]/[HA]) = 0.26, giving a ratio of about 1.82. Mix roughly 0.182 mol sodium ethanoate for every 0.100 mol ethanoic acid in 1 dm3 of water.
What is the Henderson–Hasselbalch equation and when do I use it?
The Henderson–Hasselbalch equation, pH = pKa + log([A-]/[HA]), links a buffer's pH to the pKa of the weak acid and the ratio of conjugate base to acid present. Use it whenever you're given, or can find, the moles or concentrations of both buffer components — it's the fastest route to a numerical buffer pH.
It only applies once significant amounts of both the acid and its conjugate base are present — don't use it for a weak acid solution before any of it's been neutralised. For basic buffers, use the equivalent form, pOH = pKb + log([BH+]/[B]), then convert with pH = 14 − pOH at 25°C.
Exam & Syllabus
Is buffer solutions HL only in IB Chemistry?
Yes — buffer solutions are HL-only content in the current IB Chemistry guide, sitting inside Reactivity 3.4 (pH curves). SL students study acids, bases and pH but never calculate buffer pH or design a buffer, so this topic only appears on HL Paper 1 and Paper 2.
| Aspect | SL | HL |
|---|---|---|
| Acid/base basics | Yes | Yes |
| Strong acid/base pH | Yes | Yes |
| Buffer pH calculations | No | Yes |
| Full titration pH curves | Basic | Yes, incl. buffers |
Where does buffers fit in the IB Chemistry syllabus?
Buffer solutions fall under Reactivity 3: Acids and bases, specifically sub-topic 3.4 on pH curves of acid-base titrations, in the DP Chemistry guide for first exams 2025. It follows directly from weak acid and base equilibria in Reactivity 3.3, so you need Ka, Kb and pKa comfortable before tackling buffers properly.
If your equilibrium constants are shaky, buffers will feel much harder than they are — the buffer maths is just an application of the weak acid equilibrium expression you've already met.
What common mistakes do students make with buffer solution questions?
The most frequent error I see is using Ka instead of pKa in Henderson–Hasselbalch — an easy one-mark loss. Others include forgetting the equilibrium shifts rather than going to completion when acid or base is added, and mixing up which component neutralises which addition.
Checklist before you submit a buffer answer:
- Did you take -log(Ka) to get pKa, not use Ka directly?
- Did you show the correct species reacting with added H+ versus OH-?
- Did you use moles or concentrations consistently, not mix the two?
- Does your buffer ratio make chemical sense — not negative, not wildly extreme?
How many marks are buffer questions usually worth in IB Chemistry exams?
Buffer questions typically appear as 2-5 mark parts within a larger acids-and-bases question on Paper 2, sometimes stretching to a 6-8 mark structured question combining explanation and calculation. They rarely appear in Paper 1 multiple choice since a full buffer calculation needs working space to award method marks.
Expect the mark scheme to award separate marks for identifying pKa correctly, substituting the ratio, and stating a final pH to appropriate significant figures — losing the last mark for rounding is a common, avoidable slip.
Difficulty, Grades & Revision
How hard is the buffers topic in IB Chemistry HL?
Buffers sit in the trickier half of HL acids and bases — not conceptually difficult once you see the pattern, but easy to lose marks on through calculation slips. Students who've mastered equilibrium constants and logarithms generally find it manageable; those shaky on Ka/pKa conversions struggle most.
In my experience marking mock papers, the grade gap here isn't about intelligence — it's practice with logarithms. A student comfortable rearranging pH = pKa + log(x) under time pressure usually picks up full marks; one who freezes at the log step can lose two or three marks per question.
How can my child revise buffer solutions effectively for IB Chemistry?
The most effective revision combines rebuilding the Henderson–Hasselbalch equation from memory, working through past-paper buffer calculations under timed conditions, and reviewing command terms like state, calculate and explain so your child answers exactly what's asked. Topical practice on acids and bases, checked against mark schemes, helps spot recurring mistakes early.
Three concrete steps: 1) rebuild Henderson–Hasselbalch from the equilibrium expression without looking it up. 2) work five past-paper buffer questions timed, checking each against the mark scheme for exact method marks. 3) keep a short list of every mistake made and revisit it a week later — spaced repetition catches the same one-mark slips examiners see every session.
Acidic Buffer vs Basic Buffer
| Feature | Acidic Buffer | Basic Buffer |
| Components | Weak acid + conjugate base salt | Weak base + conjugate acid salt |
| Typical pH | Below 7 | Above 7 |
| Classic IB example | Ethanoic acid / sodium ethanoate | Ammonia / ammonium chloride |
| Key equation | pH = pKa + log([A-]/[HA]) | pOH = pKb + log([BH+]/[B]) |
For more worked buffer calculations and full mark schemes, see RevisionPrep's IB Chemistry Topical Worksheets and Revision Notes on acids and bases.