AP Chemistry Unit 8 Notes: Acids, Bases, Buffers, pH/pKa, and Titrations

Brønsted–Lowry acid

A species that donates a proton (H+) in a reaction.

Brønsted–Lowry base

A species that accepts a proton (H+) in a reaction.

Conjugate acid–base pair

Two species that differ by exactly one proton; an acid forms its conjugate base after donating H+, and a base forms its conjugate acid after accepting H+.

Strong acid

An acid that essentially fully ionizes in water (e.g., HCl, HNO3), so [H3O+] is determined directly from concentration (after dilution/mixing).

Weak acid

An acid that partially ionizes in water and establishes an equilibrium (e.g., HF, CH3COOH), requiring Ka-based equilibrium calculations.

Acid dissociation constant (Ka)

Equilibrium constant for a weak acid in water: HA + H2O ⇌ A− + H3O+, with Ka = ([H3O+][A−])/[HA]. Larger Ka means a stronger weak acid.

Base dissociation constant (Kb)

Equilibrium constant for a weak base in water: B + H2O ⇌ BH+ + OH−, with Kb = ([BH+][OH−])/[B]. Larger Kb means a stronger weak base.

Ka–Kb relationship (conjugates)

For a conjugate acid–base pair HA/A− at a given temperature: Ka·Kb = Kw (so pKa + pKb = pKw).

Neutralization

An acid–base reaction where acid and base react to form water and (usually) a salt; often treated as going to completion for strong acid/strong base.

Net ionic equation (strong acid–strong base)

The essential reaction in strong acid–strong base neutralization: H3O+ + OH− → 2H2O.

Weaker acid/base pair favored

Acid–base equilibria tend to favor formation of the weaker acid and weaker base; for HA + B− ⇌ A− + HB, K = Ka(HA)/Ka(HB).

Buffer

A solution that resists pH change when small amounts of strong acid/base are added; requires a conjugate acid–base pair in comparable amounts.

Buffer components (HA/A−)

A common buffer is made from a weak acid (HA) and its conjugate base (A−), present in significant/comparable amounts (often from the acid plus its salt).

Buffer response to added strong acid

Added H3O+ is consumed by the base component: A− + H3O+ → HA + H2O, reducing the pH change.

Buffer response to added strong base

Added OH− is consumed by the acid component: HA + OH− → A− + H2O, reducing the pH change.

Henderson–Hasselbalch equation

For an HA/A− buffer: pH = pKa + log([A−]/[HA]); after adding acid/base, the ratio is often computed using moles because dilution cancels in the same total volume.

Half-equivalence point

In a weak acid–strong base titration, the point where half the initial HA has been converted to A−, so [A−] = [HA] and pH = pKa.

Buffer capacity

How much strong acid/base a buffer can absorb before pH changes significantly; increases with larger total amounts (higher concentrations) of HA and A−.

pH

A logarithmic measure of acidity based on hydronium: pH = −log([H3O+]); a 1-unit change corresponds to a 10× change in [H3O+].

pOH

A logarithmic measure of basicity: pOH = −log([OH−]); related to pH through Kw.

Ion-product constant of water (Kw)

At 25°C, Kw = [H3O+][OH−] = 1.0×10−14, linking hydronium and hydroxide concentrations.

pKw and the pH + pOH relationship

pKw = −log(Kw) = 14.00 at 25°C, so pH + pOH = 14.00 (temperature-dependent).

pKa

The log form of Ka: pKa = −log(Ka). Smaller pKa means a stronger acid (larger Ka).

Equivalence point (titration)

The point in a titration where acid and base have reacted in the exact stoichiometric ratio (moles neutralized match the reaction coefficients); not the same as endpoint.

Endpoint (titration)

The experimental indicator color-change point used to signal completion; it should be close to the equivalence point but is not defined by stoichiometry.