AP Chem Unit 8

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

The species formed when a base accepts a proton; the conjugate acid of base B is BH⁺.

Conjugate Base

The species formed when an acid donates a proton; the conjugate base of acid HA is A⁻.

Autoionization of Water

The reaction in which water acts as both acid and base: 2H₂O ⇌ H₃O⁺ + OH⁻; Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C.

pH

−log[H₃O⁺]; a measure of acidity; pH < 7 is acidic, pH = 7 is neutral, pH > 7 is basic (at 25°C).

pOH

−log[OH⁻]; a measure of basicity; pH + pOH = 14 at 25°C.

Strong Acid

An acid that completely ionizes in water; strong acids: HCl, HBr, HI, HClO₄, H₂SO₄, HNO₃; [H₃O⁺] = initial acid concentration.

Weak Acid

An acid that only partially ionizes in water; equilibrium established; Ka << 1.

Strong Base

A base that completely dissociates in water; group I and II hydroxides; [OH⁻] = initial base concentration (× 2 for group II).

Weak Base

A base that only partially ionizes in water; Kb << 1.

Acid Ionization Constant (Ka)

The equilibrium constant for the ionization of a weak acid; Ka = [H₃O⁺][A⁻]/[HA]; larger Ka = stronger acid.

Base Ionization Constant (Kb)

The equilibrium constant for the ionization of a weak base; Kb = [OH⁻][HB⁺]/[B].

pKa

−log Ka; smaller pKa = stronger acid.

pKb

−log Kb; smaller pKb = stronger base.

Ka × Kb = Kw

Relationship between a conjugate acid-base pair; pKa + pKb = pKw = 14.

Percent Ionization

The fraction of weak acid (or base) molecules that ionize in solution; increases with dilution.

ICE Table (acid-base)

Used to calculate equilibrium concentrations (and pH) of weak acid/base solutions.

Acid Strength and Molecular Structure

Electronegativity, inductive effects, and resonance stabilize the conjugate base and increase acid strength; stronger acid → weaker conjugate base.

Oxyacids

Acids containing oxygen; stronger as the number of O atoms increases or as the central atom's electronegativity increases.

Carboxylic Acid

A class of weak organic acids containing the –COOH group.

Nitrogenous Bases

Common weak bases including ammonia (NH₃) and amines.

Hydrolysis of Salts

Salts of weak acids produce basic solutions; salts of weak bases produce acidic solutions; salts of strong acid + strong base produce neutral solutions.

Titration Curve

A graph of pH vs. volume of titrant added; shape depends on whether acid and base are strong or weak.

Equivalence Point (acid-base)

The point in a titration where moles of titrant equal moles of analyte; pH depends on the identities of the acid and base.

Half-Equivalence Point

The point in a weak acid titration where [HA] = [A⁻]; at this point pH = pKa; useful for determining Ka experimentally.

Buffer Solution

A solution containing significant amounts of both a weak acid and its conjugate base; resists pH changes upon addition of small amounts of acid or base.

Henderson-Hasselbalch Equation

pH = pKa + log([A⁻]/[HA]); used to calculate the pH of a buffer solution.

Buffer Capacity

The amount of acid or base a buffer can neutralize before significant pH change occurs; increases with higher concentrations of buffer components.

Acid-Base Indicator

A weak acid or base that changes color in its protonated vs. deprotonated form; should be selected so its pKa ≈ pH at equivalence point.

Polyprotic Acid

An acid capable of donating more than one proton (e.g., H₂SO₄, H₃PO₄); has multiple Ka values; Ka₁ > Ka₂ > Ka₃.

pH and Solubility

The solubility of a salt is pH-sensitive when one of its ions is the conjugate of a weak acid or base; lower pH increases solubility of salts with basic anions.