Properties of Buffers
Explain how both weak conjugate forms consume small additions and why buffering has finite limits.
AP Chemistry Topic 8.8 · CED printed p. 156 / PDF p. 163. This teaching, practice and feedback were added for the AP edition; original author material remains separately traceable.
Before you start
- Conjugate pairs
- Quantitative proton transfer
- Major species
A useful buffer contains appreciable concentrations of both members of a weak conjugate acid–base pair. The base member consumes added hydronium; the acid member consumes added hydroxide. For HA/A−, write A− + H3O+ → HA + H2O and HA + OH− → A− + H2O. This is why a buffer resists changes in both directions. A weak acid alone normally contains too little of its conjugate base to provide the same protection against added acid. A strong acid with its extremely weak conjugate base is not a useful ordinary buffer near neutral pH.
The buffer does not make added acid or base disappear without consequence. It converts HA into A− or the reverse, changing their ratio. Because the original reservoirs are large compared with a small addition, that ratio changes modestly and pH changes modestly. Once one reservoir is largely exhausted, the protection fails. At the particle level, show a proton moving from added hydronium to A−, or from HA to hydroxide; do not draw the buffer as merely swallowing ions unchanged. Dilution can leave the pair ratio almost unchanged while shrinking its ability to absorb the same added amount per volume. These are finite-capacity solutions, not a promise of constant pH under every disturbance.
Let's work through it
- Consider 100.0 mL containing 0.0100 mol HA and 0.0100 mol A−. A small conceptual addition supplies 0.0010 mol H3O+.
- The reaction A− + H3O+ → HA + H2O consumes most of the added hydronium. HA increases and A− decreases; both remain appreciable.
- The buffer’s pH therefore moves downward somewhat rather than as dramatically as an unbuffered solution. No numerical pH-change calculation is needed for this AP claim.
- If instead the added acid amount exceeded the available A−, excess strong acid would remain and the original buffer protection would fail.
Make the picture move
Interactive concept model
The species change across a titration curve.
Predict first. Visit 0, 12.5, 25.0, and 37.5 mL. Identify the species controlling pH at each point. Is equivalence necessarily pH 7?
HA + OH⁻ → A⁻ + H₂O; Kₐ = [H⁺][A⁻]/[HA]; K_w = [H⁺][OH⁻]
- Ideal weak-acid titration
- Labeled selected or reference state
On a small screen, scroll inside the graph or diagram to read its labels. The values and explanation also appear below.
At 12.5 mL, pH = 4.74. Half-equivalence: pH ≈ pKₐ. [HA] = 0.0333 M and [A⁻] = 0.0334 M.
Read the data and comparisons
| Added base (mL) | pH | [HA] (mol L⁻¹) | [A⁻] (mol L⁻¹) |
|---|---|---|---|
| 0.0 | 2.87 | 0.0987 | 0.00134 |
| 12.5 | 4.74 | 0.0333 | 0.0334 |
| 25.0 | 8.72 | 5.24 × 10^-6 | 0.05 |
| 37.5 | 12.30 | 1.10 × 10^-9 | 0.04 |
| 50.0 | 12.52 | 5.50 × 10^-10 | 0.0333 |
Macroscopic / data. The pH curve is a computed data representation, with a changing total liquid volume.
Particle view. HA, A⁻, H⁺, OH⁻, and spectator Na⁺ satisfy matter and charge conservation. Equivalence does not remove the hydrolysis equilibrium.
Symbolic view. Total acid-species concentration = 0.0667 M; [H⁺] = 1.82 × 10^-5 M; [OH⁻] = 5.50 × 10^-10 M.
Why it works. The initial 25.0 mL of 0.100 M HA contains 0.00250 mol of acid. Equivalence takes 25.0 mL of 0.100 M base regardless of pKₐ. At half-equivalence, comparable HA and A⁻ give pH approximately pKₐ. Within the buffer region, pH = pKₐ + log₁₀([A⁻]/[HA]) follows by rearranging Kₐ; estimating this ratio from neutralization requires appreciable amounts of both species. At equivalence A⁻ reacts with water, so a weak-acid titration is basic. Beyond equivalence excess OH⁻ dominates.
Conditions and limits. Monoprotic hypothetical HA and a fully dissociated strong base with Na⁺ as spectator; 25 °C, K_w = 1.0 × 10⁻¹⁴, ideal dilute activities approximated by concentrations divided by a 1 mol L⁻¹ standard state, additive volumes, no other equilibria. The curve solves acid mass balance, charge balance, Kₐ, and K_w together at every volume; it does not misuse Henderson–Hasselbalch at endpoints. The numerical solver is an implementation detail, not required AP work. pH is displayed to two decimal places.
Try it from memory
Pause & retrieve
Try it before you peek.
Write your answer and the reason from memory. Then compare with the worked answer. Your self-assessment is a study decision, not an automatically graded score.
Question 1
Which member removes added hydroxide?
See every recall question with its answer
These references preserve each prompt and its answer for direct evidence links. Try the interactive questions above before opening the worked reasoning.
Retrieval 1
Which member removes added hydroxide?
Worked answer and reason
HA, by donating a proton and becoming A−.
The conjugate acid reacts with added base.
Retrieval 2
Why does a buffer eventually fail?
Worked answer and reason
One component can be depleted relative to the added acid/base amount.
Its neutralizing reservoir is finite.
Catch the tempting mistake
- A buffer keeps pH exactly constant.
- Its component ratio changes on addition, so pH changes too, usually less than without the buffer.
Put it to work
Give the chemical reason as well as the calculation. Worked solutions are on a separate page.
- Unit 8 problem 12 · ap application · particle
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For the teacher
- Require the two neutralization equations and a before/after particle explanation. Quantitative buffer pH-change calculations are outside AP assessment.
Source objectives and scope
Source edition: ap-chemistry-ced-fall2024-local-b5dfe867. See this topic in the course outline →
- 8.8.A
- Explain the relationship between the ability of a buffer to stabilize pH and the reactions that occur when an acid or a base is added to a buffered solution.
- 8.8.A.1
- A buffer solution contains a large concentration of both members in a conjugate acid-base pair. The conjugate acid reacts with added base and the conjugate base reacts with added acid. These reactions are responsible for the ability of a buffer to stabilize pH.
Science references for this teaching
- AP Chemistry Course and Exam Description, effective Fall 2024 — Required topic-numbered learning objectives and essential knowledge, all continuations and AP exclusions; supplied PDF SHA256 b5dfe8677ef3d88c613865d2e2a3e8d6125d652e2b24c71ef1e8ce4e011094f0.
- OpenStax Chemistry 2e §14.6, Buffers — Conjugate-pair buffer reactions, composition ratio, and capacity.
Glossary and deeper reading · 6 refreshers
These editorial refreshers supplement the reading. Definitions and examples are added guidance; original source wording and recorded answers remain in their original context.
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Conjugate acid–base pair
A conjugate acid–base pair consists of species differing by one proton. Losing a proton forms an acid's conjugate base, while gaining one forms a base's conjugate acid. For example, ammonium and ammonia form a conjugate acid–base pair.
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Brønsted acid and base
A Brønsted acid donates a proton, and a Brønsted base accepts a proton. These roles describe a particular reaction, so some substances can act in either role. For example, water accepts a proton from hydrogen chloride but donates one to ammonia.
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Hydronium
Hydronium is the H₃O⁺ ion used in the introductory model of proton behavior in water. Aqueous H⁺ notation is shorthand for a solvated proton rather than a bare proton moving independently in the liquid. For example, a Brønsted acid can transfer a proton to water and form hydronium.
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Weak acid or base
A weak acid or base reacts only partially with water and establishes an equilibrium. Its equilibrium constant describes strength, while analytical concentration describes how much was prepared per volume. For example, acetic acid leaves both acid molecules and its conjugate base in water.
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pH
pH is the negative base-ten logarithm of hydronium activity. In suitable dilute models, activity is approximated by molar concentration relative to the standard concentration. For example, lowering hydronium concentration by a factor of ten raises pH by one unit.
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Ion
An ion is an atom or group of atoms with a net electric charge. Ordinary ion formation changes electron count without changing the element's proton count. For example, a sodium atom losing one electron becomes a sodium ion with charge +1.