Buffer Solutions

8.8 Properties of Buffers

Add acid or base and watch buffered vs. unbuffered pH curves split apart. Animated particles show the HA/A⁻ ratio shift as Henderson-Hasselbalch recalculates.

Henderson-HasselbalchBuffer CapacitypH Curves5 Acids

Topic 8.8

Properties of Buffers

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.

A buffer solution contains a large concentration of both members of a conjugate acid–base pair. The conjugate acid reacts with added base, and the conjugate base reacts with added acid. Those two reactions are what allow a buffer to stabilize pH.

The two neutralization reactions:

Add acid: A⁻ + H₃O⁺ → HA + H₂O
Add base: HA + OH⁻ → A⁻ + H₂O

Notice that neither added species survives as free H₃O⁺ or OH⁻ — each is converted into a member of the pair already present. All that changes is the ratio, and since pH depends on the log of the ratio, a modest change in the ratio produces a very small change in pH.

How to make a buffer. Three routes, all producing the same thing:

Mix a weak acid with a soluble salt of its conjugate base (acetic acid + sodium acetate).
Partially neutralize a weak acid with a strong base — stop before the equivalence point.
Partially neutralize a weak base with a strong acid.

What is not a buffer: a strong acid with its conjugate base (Cl⁻ is far too weak a base to react with added acid), or a weak acid alone with no conjugate base present.

Connection to the common-ion effect (7.12). Adding acetate to acetic acid suppresses the acid's ionization exactly as adding chloride suppresses AgCl's dissolution — same principle, different equilibrium.

Biological relevance: blood is buffered by the H₂CO₃/HCO₃⁻ pair at pH ≈ 7.4, holding within about ±0.05 units despite continuous metabolic acid production.

Key points

A buffer needs meaningful concentrations of both members of a conjugate pair.
Added acid is consumed by A⁻; added base is consumed by HA.
Only the ratio changes, and pH depends on the log of the ratio — hence the small response.
A strong acid and its conjugate base cannot buffer.

Equations

$A^{-} + H_{3} O^{+} \to HA + H_{2} O$
not on the sheetBuffer response to added acid.
$HA + O H^{-} \to A^{-} + H_{2} O$
not on the sheetBuffer response to added base.

Common mistakes

Calling a strong acid + its salt a buffer. HCl/NaCl does not buffer.
Thinking a buffer holds pH exactly constant. It resists change; it does not prevent it.
Forgetting a buffer works in both directions. It absorbs added acid and added base.
Assuming any weak acid solution is a buffer. The conjugate base must also be present in quantity.

Worked example

A buffer is prepared from 0.50 M HF and 0.50 M NaF. Write the equations showing how it responds to added HCl and to added NaOH, and explain why the pH changes only slightly.

Response to added HCl (which supplies H₃O⁺):
F⁻(aq) + H₃O⁺(aq) → HF(aq) + H₂O(l)

The fluoride ion, present in large concentration, consumes the added hydronium ion and converts it into HF. Virtually none of the added H₃O⁺ remains free in solution.

Response to added NaOH (which supplies OH⁻):
HF(aq) + OH⁻(aq) → F⁻(aq) + H₂O(l)

The hydrofluoric acid, also present in large concentration, consumes the added hydroxide and converts it into F⁻.

Why pH barely moves. The pH of the buffer is set by pH = pK_a + log([F⁻]/[HF]). Adding a small amount of acid or base converts a small quantity of one member of the pair into the other, changing the ratio only slightly. Because pH depends on the logarithm of that ratio, even a change from 1.00 to 1.20 alters the pH by only 0.08 units. Contrast this with unbuffered water, where adding the same amount of strong acid could drop the pH by several units, because there is no reservoir species to consume it.

Full notes for topic 8.8 →