Buffers and Acid-Base Titrations Study Pack
Kibin's free study pack on Buffers and Acid-Base Titrations includes a 7-section study guide, 25 quiz questions, 30 flashcards, and 5 open-ended Explain review questions. Sign up free to track your progress toward mastery, plus upload your own notes and recordings to create personalized study packs organized by course.
Last updated May 27, 2026
Buffers and Acid-Base Titrations Study Guide
Master the mechanics of buffers and acid-base titrations, from applying the Henderson-Hasselbalch equation and maximizing buffer capacity to identifying equivalence points, half-equivalence points, and choosing the right indicator for any titration curve.
Key Takeaways
- •A buffer is a solution that resists changes in pH when small amounts of strong acid or strong base are added, and it functions by containing both a weak acid and its conjugate base (or a weak base and its conjugate acid) in comparable concentrations.
- •The Henderson-Hasselbalch equation, pH = pKa + log([A⁻]/[HA]), allows calculation of buffer pH from the ratio of conjugate base to weak acid concentrations.
- •A buffer's capacity — its ability to neutralize added acid or base without significant pH change — is greatest when the concentrations of the weak acid and conjugate base are equal, meaning pH ≈ pKa.
- •In an acid-base titration, a solution of known concentration (the titrant) is added incrementally to a solution of unknown concentration (the analyte) until the reaction reaches the equivalence point, where moles of acid exactly equal moles of base.
- •The equivalence point pH depends on the nature of the acid-base pair: strong acid–strong base titrations reach equivalence at pH 7, while weak acid–strong base titrations reach equivalence at pH > 7 due to hydrolysis of the conjugate base.
- •An indicator signals the endpoint of a titration by changing color when the solution pH crosses a specific range, and it should be selected so that its transition range overlaps with the equivalence point pH.
- •The half-equivalence point in a weak acid–strong base titration is the stage at which exactly half the weak acid has been neutralized; at this point, [HA] = [A⁻] and pH equals the pKa of the weak acid.
How Buffers Work: Components and Mechanism
A buffer resists pH change because it contains two chemical species — one capable of neutralizing added acid and one capable of neutralizing added base — that exist in equilibrium with each other.
Two Types of Buffer Systems
- •A weak acid–conjugate base buffer (e.g., acetic acid / acetate ion, CH₃COOH / CH₃COO⁻) resists acidification because the conjugate base consumes added H⁺, and resists basification because the weak acid donates H⁺ to neutralize added OH⁻.
- •A weak base–conjugate acid buffer (e.g., ammonia / ammonium ion, NH₃ / NH₄⁺) works by the same logic in reverse: NH₃ absorbs added H⁺ and NH₄⁺ releases H⁺ when base is added.
- •Both types require that the weak acid (or conjugate acid) and its paired base be present in significant, comparable amounts — typically within a factor of 10 of each other.
Molecular-Level Neutralization Reactions
- •When a strong acid such as HCl is added to a weak acid–conjugate base buffer, the H⁺ ions react with CH₃COO⁻ to form CH₃COOH, consuming the added acid before it can shift the pH.
- •When a strong base such as NaOH is added, OH⁻ reacts with CH₃COOH to produce CH₃COO⁻ and water, again preventing a large pH spike.
- •These neutralization reactions are not perfectly efficient — large additions of strong acid or base eventually overwhelm the buffer by depleting one of its components.
The Henderson-Hasselbalch Equation and Buffer pH
The quantitative relationship between buffer composition and pH is captured by the Henderson-Hasselbalch equation, which is derived directly from the equilibrium expression for the dissociation of a weak acid.
Derivation from the Weak Acid Equilibrium
- •Starting from Ka = [H⁺][A⁻]/[HA], solving for [H⁺] gives [H⁺] = Ka × ([HA]/[A⁻]). Taking the negative logarithm of both sides yields pH = pKa + log([A⁻]/[HA]).
- •This form shows that buffer pH is anchored to the pKa of the weak acid and then shifted up or down depending on the ratio of conjugate base to weak acid.
Interpreting the Ratio [A⁻]/[HA]
- •When [A⁻] = [HA], the log term equals zero and pH = pKa exactly — this is the point of maximum buffer capacity.
- •When [A⁻] > [HA], the buffer is more basic than the pKa; when [HA] > [A⁻], the buffer is more acidic.
- •Practical buffers are effective within roughly one pH unit on either side of the pKa, corresponding to concentration ratios between 0.1 and 10.
Selecting the Right Weak Acid for a Target pH
- •To prepare a buffer at a desired pH, choose a weak acid whose pKa is within one unit of that target, then adjust the [A⁻]/[HA] ratio to fine-tune the pH.
- •For example, to prepare a buffer at pH 4.7, acetic acid (pKa ≈ 4.74) is an appropriate choice; adjusting the acetate-to-acetic-acid ratio brings the pH precisely to target.
Buffer Capacity and Limitations
Buffer capacity describes how much strong acid or strong base a buffer can absorb before its pH changes significantly, and it depends on both the absolute concentrations of the buffer components and their ratio.
Factors That Determine Buffer Capacity
- •Higher total concentrations of weak acid and conjugate base increase buffer capacity, because more moles of neutralizing agent are available.
- •A 1:1 ratio of [A⁻] to [HA] maximizes capacity because neither component is close to exhaustion; a heavily lopsided ratio depletes one component quickly.
- •Buffer capacity is asymmetric when the ratio is not 1:1 — a buffer with excess weak acid absorbs more base than acid before failing.
When Buffers Fail
- •Adding strong acid beyond what the conjugate base can absorb converts essentially all A⁻ to HA, removing the species that neutralizes acid; pH then drops sharply.
- •Similarly, adding excess strong base converts all HA to A⁻, removing the acid component; pH rises sharply.
- •These failure points define the practical boundaries of a buffer's working range.
Acid-Base Titration: Setup, Progress, and the Equivalence Point
An acid-base titration is a quantitative technique in which a solution of precisely known concentration — the titrant — is delivered from a burette into a measured volume of the analyte until the neutralization reaction is complete.
Core Components of a Titration
- •The analyte is the solution whose concentration is unknown; the titrant is the standardized solution added in controlled volumes using a burette.
- •A primary standard (e.g., potassium hydrogen phthalate, KHP, for base standardization) is a highly pure compound used to establish the exact concentration of the titrant before use.
- •An indicator or pH meter detects when the reaction has reached completion.
The Equivalence Point Defined
- •At the equivalence point, the moles of H⁺ supplied by the acid exactly equal the moles of OH⁻ supplied by the base — no excess of either remains.
- •For a monoprotic acid–monobase titration, this condition is expressed as: moles of acid = moles of base, or Ca × Va = Cb × Vb.
- •The equivalence point is a calculated or theoretically defined position; the endpoint is the experimentally observed color change of the indicator, which ideally coincides with it.
pH at the Equivalence Point: Strong vs. Weak Systems
- •A strong acid–strong base titration (e.g., HCl titrated with NaOH) produces only water and a neutral salt at equivalence; pH = 7 at 25°C.
- •A weak acid–strong base titration (e.g., acetic acid titrated with NaOH) produces the conjugate base of the weak acid (CH₃COO⁻) at equivalence; this conjugate base hydrolyzes water to give OH⁻, raising the equivalence point pH above 7.
- •A weak base–strong acid titration produces the conjugate acid at equivalence, which lowers the equivalence point pH below 7.
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What is the pH of a buffer solution when the concentrations of the weak acid and its conjugate base are exactly equal?
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How Buffers Work
Explain how a buffer resists changes in pH in your own words. What components must be present, and what happens at the molecular level when a small amount of strong acid or strong base is added to a buffer?
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