AP Chemistry Study Guide

Acids and bases is the single most tested unit on AP Chemistry. Expect 2–4 free-response points and multiple MCQs directly from this unit.

The pH Scale

At 25°C, a neutral solution has pH = 7. Below 7 is acidic; above 7 is basic.

Strong vs. Weak Acids

Strong acids dissociate 100% in water. The 6 to memorize: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄.

Weak acids partially dissociate. A larger Ka means a stronger acid. Use an ICE table to find [H⁺] from Ka.

ICE Table Example

For 0.10 M acetic acid (Ka = 1.8 × 10⁻⁵):

         CH₃COOH  ⇌  H⁺  +  CH₃COO⁻
Initial:   0.10        0       0
Change:    −x         +x      +x
Equil:   0.10−x       x       x

Ka = x² / (0.10 − x) ≈ x² / 0.10
x = [H⁺] = √(1.8×10⁻⁵ × 0.10) ≈ 1.34×10⁻³ M
pH = −log(1.34×10⁻³) ≈ 2.87

Buffers and Henderson-Hasselbalch

A buffer contains a weak acid (HA) and its conjugate base (A⁻) — typically from a salt. When acid is added, A⁻ absorbs it. When base is added, HA neutralizes it.

Key buffer facts for the exam:

• Buffer capacity is highest when pH = pKa (equal concentrations of HA and A⁻).
• Adding a strong acid to a buffer: convert moles, recalculate ratio, plug into H-H.
• A buffer with high concentrations resists larger pH changes.

Titrations

At the equivalence point of a weak acid–strong base titration, the solution is basic (the conjugate base hydrolyzes). For strong acid–strong base, it's exactly neutral (pH = 7).

The half-equivalence point is where pH = pKa — a great shortcut on FRQs.

Chemical equilibrium occurs when the forward and reverse reaction rates are equal. The position of equilibrium is described by the equilibrium constant K.

Writing K Expressions

For the reaction aA + bB ⇌ cC + dD (pure solids and liquids omitted):

The Reaction Quotient Q

Q has the same form as K but uses current (non-equilibrium) concentrations.

• Q < K: Reaction shifts right (toward products).
• Q > K: Reaction shifts left (toward reactants).
• Q = K: At equilibrium.

Le Chatelier's Principle

When a system at equilibrium is disturbed, it shifts to partially counteract the change.

Disturbance	Shift Direction
Add reactant	Right (toward products)
Remove product	Right
Increase pressure (gas)	Toward fewer moles of gas
Increase temperature	Toward endothermic direction
Add inert gas (constant V)	No shift

Solubility Equilibria (Ksp)

For a sparingly soluble salt CaF₂ ⇌ Ca²⁺ + 2F⁻:

The common-ion effect: adding a common ion decreases solubility. Adding NaF to a CaF₂ solution will make less CaF₂ dissolve.

Enthalpy (ΔH)

Enthalpy measures heat flow at constant pressure. Exothermic reactions release heat (ΔH < 0); endothermic reactions absorb heat (ΔH > 0).

Hess's Law: ΔH for a reaction equals the sum of ΔH values for steps that add up to the overall reaction. Flip a reaction → flip the sign of ΔH.

Entropy (ΔS)

Entropy measures disorder. Processes that increase entropy (ΔS > 0):

• Solids → liquids → gases
• More moles of gas formed than consumed
• Dissolution of ionic solids (usually)
• Mixing of substances

Gibbs Free Energy (ΔG)

A reaction is spontaneous when ΔG < 0.

ΔH	ΔS	Spontaneous?
− (exo)	+ (more disorder)	Always
+ (endo)	− (less disorder)	Never
− (exo)	− (less disorder)	Low T only
+ (endo)	+ (more disorder)	High T only

Rate Laws

The rate law expresses how reaction rate depends on concentration. Exponents m and n are determined experimentally — never from stoichiometry. Overall order = m + n.

Integrated Rate Laws

Order	Integrated Rate Law	Half-Life
0	[A] = [A]₀ − kt	t½ = [A]₀ / 2k
1	ln[A] = ln[A]₀ − kt	t½ = 0.693 / k (constant)
2	1/[A] = 1/[A]₀ + kt	t½ = 1 / (k[A]₀)

Arrhenius Equation

Rate constants increase with temperature. Ea is the activation energy; a catalyst lowers Ea without being consumed.

Intermolecular forces (IMFs) act between molecules — they are weaker than intramolecular (covalent/ionic) bonds but control physical properties.

Types of IMFs (weakest to strongest)

1. London Dispersion Forces (LDF) — present in ALL molecules. Caused by temporary dipoles from electron motion. Stronger for larger, heavier, more polarizable molecules.
2. Dipole-Dipole — between polar molecules (permanent dipoles). Present when molecules have a net dipole moment.
3. Hydrogen Bonding — special dipole-dipole between H bonded to N, O, or F and a lone pair on another N, O, or F. Much stronger than regular dipole-dipole.
4. Ion-Dipole — between an ion and a polar molecule. Strongest of the intermolecular forces; important in dissolving ionic compounds in water.

IMFs and Physical Properties

Stronger IMFs → higher boiling point, higher viscosity, lower vapor pressure.

• Water has an anomalously high boiling point (100°C) for its mass because of extensive hydrogen bonding.
• Branched molecules have lower boiling points than straight-chain isomers (less surface contact → weaker LDF).