AP Chemistry Unit 8 Buffers: How Solutions Resist pH Change

Buffer

A solution that resists changes in pH when small amounts of strong acid or strong base are added (pH changes much less than in pure water).

Conjugate acid–base pair

Two species that differ by one proton (H+), such as HA/A− or B/BH+, required in appreciable amounts to make a buffer.

Weak acid (HA)

An acid that only partially dissociates in water; in a buffer it neutralizes added OH−.

Conjugate base (A−)

The base formed when a weak acid loses H+; in a buffer it neutralizes added H3O+ (often supplied by a soluble salt like NaA).

Acid buffer (HA/A−)

A buffer made from a weak acid HA and its conjugate base A−, providing “two-way” neutralization of added acid or base.

Base buffer (B/BH+)

A buffer made from a weak base B and its conjugate acid BH+, used to resist pH changes around a basic pH.

Why strong acids/bases don’t make buffers

Strong acids/bases essentially fully react with water, leaving no meaningful equilibrium “reserve” of conjugate forms to respond gently to added acid/base.

Neutralization of added strong acid (acid buffer)

A− + H3O+ → HA + H2O; the conjugate base consumes added acid, converting it to weak acid.

Neutralization of added strong base (acid buffer)

HA + OH− → A− + H2O; the weak acid consumes added base, limiting pH increase.

Buffer pH dependence on ratio

Buffer pH depends primarily on the relative amounts (ratio) of conjugate base to weak acid (or B to BH+), not just the presence of an acid/base.

pKa

−log Ka; a measure of acid strength (smaller pKa means stronger acid) and a key reference point for buffer effectiveness.

Buffer region (effective buffer range)

The pH range where a buffer works well, typically pH ≈ pKa ± 1 for an HA/A− buffer.

Tenfold ratio change rule

A 1-unit change in pH corresponds to a tenfold change in the ratio [A−]/[HA] (because of the base-10 logarithm).

Acid dissociation equilibrium (for a buffer)

HA + H2O ⇌ H3O+ + A−; the equilibrium shifts as HA and A− amounts change during buffering.

Acid dissociation constant (Ka) expression

Ka = ([H3O+][A−])/[HA]; relates buffer component concentrations to hydronium concentration.

Henderson–Hasselbalch equation

pH = pKa + log([A−]/[HA]); a rearrangement of the Ka expression used to compute buffer pH from the conjugate pair ratio.

Meaning of “log” in Henderson–Hasselbalch

log is base-10; if [A−]/[HA] increases by a factor of 10, the log term increases by 1 and pH increases by 1.

Stoichiometry-first buffer method

After adding strong acid/base, first do neutralization stoichiometry to update moles of HA and A−, then apply Henderson–Hasselbalch.

Using moles instead of molarity (ratio shortcut)

Because Henderson–Hasselbalch uses a ratio, you can often use nA−/nHA instead of concentrations if both are in the same solution volume (or volume change is negligible/handled consistently).

Special case: [A−] = [HA]

When conjugate base and weak acid are equal, log(1)=0 and pH = pKa (often a point of maximum buffering symmetry).

Buffer capacity

How much strong acid or strong base a buffer can absorb before its pH changes dramatically; depends on total amounts and how balanced the pair is.

What buffer capacity depends on

(1) Total concentration/moles of HA and A− (more = higher capacity) and (2) how close [A−] and [HA] are (capacity greatest when comparable, often near pH ≈ pKa).

Effect of dilution on buffers

Diluting both components equally keeps [A−]/[HA] about the same so pH changes little, but capacity decreases because absolute concentrations/moles are lower.

Identifying a buffer (example: acetic acid/acetate)

A mixture of a weak acid and its conjugate base (e.g., HC2H3O2 and C2H3O2− supplied by NaC2H3O2) forms a buffer because both conjugate forms are present.

Non-buffer example (HCl/NaCl)

Mixing a strong acid with its salt (HCl and NaCl) does not form a buffer because HCl is strong (no reserve) and Cl− is an extremely weak base.