Complete AP Chemistry Review Packet

What This AP Chemistry Review Packet Actually Covers

You're here because you need to pass AP Chemistry. Not "understand it deeply" or "develop a love for chemistry"—pass it. This packet cuts through the fluff and gives you exactly what you need: the concepts, equations, and test strategies that'll get you a 4 or 5.

Everything below is organized by topic. Skip around. Focus on what you don't know. If a section looks familiar, skim it. If it doesn't, memorize it.

Atomic Structure & The Math Behind It

The exam tests atomic structure relentlessly. Know this cold or lose easy points.

The Four Quantum Numbers

Every electron in an atom is described by four quantum numbers:

• Principal (n) — energy level, 1-7. Bigger = higher energy.
• Angular momentum (l) — orbital shape. 0 = s, 1 = p, 2 = d, 3 = f.
• Magnetic (m_l) — orientation in space. Ranges from -l to +l.
• Spin (m_s) — either +½ or -½. Two electrons per orbital, opposite spins.

The Aufbau principle, Hund's rule, and the Pauli exclusion principle determine electron configurations. You must be able to write them for any element and identify exceptions (Cr, Cu, Mo, Ag, etc. all have irregular configurations).

Photoelectron Spectroscopy (PES)

PES graphs show ionization energy on the y-axis and binding energy on the x-axis. Higher peaks = more electrons at that energy level. The spacing between peaks tells you about sublevels.

Low binding energy electrons are easiest to remove. High binding energy electrons are core electrons, close to the nucleus.

Periodicity: Trends That Actually Matter

These trends govern almost every other concept in chemistry. Get them wrong, and everything else falls apart.

• Atomic radius — increases down and left. More shielding, bigger shell.
• Ionization energy — increases up and right. Harder to remove electrons from smaller atoms with less shielding.
• Electronegativity — increases up and right. Fluorine is the most electronegative element at 3.98.
• Electron affinity — generally increases up and right, but exceptions exist. More negative = more energy released when adding an electron.
• Metallic character — increases down and left. Opposite of everything else.

Ion size matters: Cations are smaller than their parent atoms. Anions are larger. Isoelectronic species (same electron count) get smaller as nuclear charge increases.

Chemical Bonding: Ionic, Covalent, and Everything Between

Types of Bonds

Ionic bonds form between metals and nonmetals. Electrons transfer completely. The resulting compound has a lattice structure with high melting points.

Covalent bonds form between nonmetals. Electrons share. Bond polarity depends on electronegativity difference:

• 0-0.4 = nonpolar covalent
• 0.4-1.7 = polar covalent
• >1.7 = ionic

Metallic bonds occur in metals. Delocalized electrons create a "sea" that conducts electricity and heat.

Lewis Structures

Draw them correctly or suffer the consequences. Steps:

1. Count total valence electrons
2. Connect atoms with single bonds
3. Complete octets (except H, which gets 2)
4. Place remaining electrons on central atom
5. Form double/triple bonds if needed to satisfy octets
6. Check formal charges (they should be zero or close to it)

Resonance structures happen when multiple valid Lewis structures exist. The real structure is a hybrid—delocalized electrons spread across the molecule.

VSEPR & Molecular Geometry

Molecular shape determines function. Here's what you need:

Lone pairs distort bond angles. Bond angles decrease when lone pairs are present because lone pair-lone pair repulsion > lone pair-bond pair > bond pair-bond pair.

Intermolecular Forces (IMFs)

IMFs determine boiling point, solubility, and phase. Know them in order of strength:

• Ion-dipole — strongest, between ions and polar molecules
• Hydrogen bonding — H bonded to F, O, or N. Stronger than regular dipole-dipole.
• Dipole-dipole — between polar molecules
• London dispersion forces — weakest, between all molecules. Strength increases with molecular weight and surface area.

Hydrogen bonding explains why water boils at 100°C instead of -80°C. It's also why HF has a higher boiling point than HCl.

States of Matter: Gases, Liquids, Solids, and Solutions

Gas Laws

The ideal gas law is your bread and butter:

PV = nRT

P = pressure (atm), V = volume (L), n = moles, R = 0.0821 L·atm/mol·K, T = temperature (K)

Use this for:

• Finding unknown values when three variables are known
• Calculating density: d = PM/RT
• Determining molar mass: M = dRT/P

Know these relationships too:

• Boyle's Law: P₁V₁ = P₂V₂ (constant T)
• Charles's Law: V₁/T₁ = V₂/T₂ (constant P)
• Avogadro's Law: V/n = constant (constant P, T)
• Dalton's Law: P_total = P₁ + P₂ + P₃...
• Graham's Law: rate₁/rate₂ = √(M₂/M₁)

Real gas behavior deviates from ideal when pressure is high and temperature is low. Particles are close together and attractions/repulsions matter. The van der Waals equation corrects for this.

Phase Diagrams

Phase diagrams show solid, liquid, and gas phases. Three regions. Three boundary lines. One triple point (all three phases in equilibrium). One critical point (above which there is no distinct liquid/gas phase).

Slope of solid-liquid line matters. Water's is negative (ice is less dense than liquid water). Most substances have positive slopes.

Solutions & Solubility

Molarity (M) = moles solute / liters solution

Molality (m) = moles solute / kg solvent

Dilution: M₁V₁ = M₂V₂

Colligative properties depend on particle count, not identity:

• Boiling point elevation: ΔT_b = K_b · m · i
• Freezing point depression: ΔT_f = K_f · m · i
• Osmotic pressure: Π = iMRT

The van't Hoff factor (i) accounts for dissociation. NaCl in water produces 2 particles. CaCl₂ produces 3.

Stoichiometry: The Math That Makes Chemistry Work

Stoichiometry is about ratios. Master these conversions or fail.

Balancing Equations

Count atoms on both sides. Adjust coefficients only—never subscripts. Then verify.

Mole Calculations

The mole is a bridge. Master this chain:

grams ↔ moles ↔ particles (atoms, molecules, formula units) ↔ liters (at STP)

1 mole = 6.022 × 10²³ particles = 22.4 L at STP (for gases)

Limiting Reagent & Percent Yield

Find the limiting reagent by comparing mole ratios to stoichiometric ratios. The limiting reagent runs out first and determines maximum product yield.

Percent yield = (actual yield / theoretical yield) × 100

Theoretical yield is what math says should happen. Actual yield is what you got in the lab. The difference is always losses and errors.

Titrations

Titrations measure concentration. The key equation:

M_aV_a/n_a = M_bV_b/n_b

At equivalence point, moles acid = moles base. Know how to calculate pH at various points in a titration, especially the half-equivalence point (where pH = pK_a for weak acids).

Thermochemistry: Energy Changes

Key Equations

q = mcΔT

q = heat (J), m = mass (g), c = specific heat capacity (J/g·°C), ΔT = temperature change

ΔH = ΔU + Δ(PV)

For most reactions at constant pressure: ΔH ≈ ΔU

Enthalpy Calculations

Hess's Law: ΔH for a reaction is the sum of ΔH values for individual steps. Add reactions, add enthalpies.

Born-Haber cycles calculate lattice energy through formation enthalpy, atomization, ionization, electron affinity, and sublimation. It's Hess's Law applied to ionic compounds.

Bond enthalpy: ΔH_rxn = (bonds broken) - (bonds formed)

Calorimetry

In calorimetry, heat lost = heat gained. If a reaction happens in water, the water absorbs the heat. Measure temperature change, calculate q, work backwards if needed.

Coffee cup calorimeter = constant pressure (for solutions). Bomb calorimeter = constant volume (for combustion).

Kinetics: Reaction Rates

Kinetics answers: how fast does this happen? And why?

Rate Laws

Rate = k[A]^m[B]ⁿ

k = rate constant. Exponents are determined experimentally—not from stoichiometry (unless the reaction is elementary).

Zero order: rate is constant, independent of concentration. Half-life is constant.

First order: rate depends on [A] only. Half-life is constant: t_1/2 = 0.693/k. [A]_t = [A]₀e^-kt

Second order: rate depends on [A]². Half-life increases as [A] decreases.

Integrated Rate Laws & Graphs

Use graphs to determine order:

• Plot [A] vs. time → straight line = zero order
• Plot ln[A] vs. time → straight line = first order
• Plot 1/[A] vs. time → straight line = second order

Slope of the linear plot = -k (first order), = k (second order).

Activation Energy & Arrhenius

k = Ae^-Ea/RT

Higher temperature = faster reaction. Higher activation energy = slower reaction.

Find E_a from two rate constants at different temperatures:

ln(k₂/k₁) = -E_a/R (1/T₂ - 1/T₁)

Reaction Mechanisms

Mechanisms are step-by-step pathways. The slow step (rate-determining step) sets the rate law. Intermediates appear in steps but cancel out in the overall reaction.

Collision Theory

For a reaction: molecules must collide with sufficient energy (≥ E_a) and correct orientation. That's it.

Equilibrium: When Opposites Match

Equilibrium means forward rate = reverse rate. Concentrations stop changing, but reactions don't stop—they continue at equal rates.

The Equilibrium Constant (K)

K = [products]^coeffs / [reactants]^coeffs

Solids and liquids are omitted from the expression. Only gases and aqueous species count.

• K >> 1: products favored
• K << 1: reactants favored
• K ≈ 1: significant amounts of both

K_p uses partial pressures. K_c uses concentrations. K_p = K_c(RT)^Δn

Le Châtelier's Principle

When you disturb equilibrium, it shifts to counteract the change.

• Add reactant → shifts right (toward products)
• Remove product → shifts right
• Increase pressure → shifts toward fewer moles of gas
• Increase temperature → shifts in endothermic direction
• Add catalyst → no shift, but equilibrium reached faster

Reaction Quotient (Q)

Q has the same formula as K. Compare Q to K:

• Q < K: reaction proceeds forward (toward products)
• Q > K: reaction proceeds reverse (toward reactants)
• Q = K: at equilibrium

Equilibrium Calculations

Set up ICE tables (Initial, Change, Equilibrium). Solve for unknowns. Plug into K expression. Calculate.

When K is very small (<10^-3), you can approximate. When K is large (>10³), the reaction essentially goes to completion.

Acids and Bases: The pH Grind

This section is worth roughly 25% of the free response. Know it cold.

pH, pOH, and [H⁺]

pH = -log[H⁺]

pOH = -log[OH^-]

pH + pOH = 14 (at 25°C)

Strong acid: [H⁺] = original concentration. Strong base: [OH^-] = original concentration.

Weak Acids & Bases

K_a = [H⁺][A^-] / [HA]

K_b = [OH^-][BH⁺] / [B]

K_a × K_b = K_w = 10^-14

For weak acids: [H⁺] ≈ √(K_a × [HA])

For weak bases: [OH^-] ≈ √(K_b × [B])

Percent ionization = (concentration ionized / initial concentration) × 100. It increases as concentration decreases for weak acids and bases.

Buffers

Buffers resist pH change. They're mixtures of weak acid and conjugate base (or weak base and conjugate acid).

Henderson-Hasselbalch:

pH = pK_a + log([A^-]/[HA])

Buffers work best when [A^-] ≈ [HA] (pH ≈ pK_a). Capacity matters—more moles of conjugate acid/base = more resistance to pH change.

Titration Curves

Know the shape of each curve:

• Strong acid + strong base: smooth curve, pH = 7 at equivalence
• Weak acid + strong base: pH starts higher, buffer region visible, pH > 7 at equivalence
• Weak base + strong acid: pH starts lower, buffer region visible, pH < 7 at equivalence

Half-equivalence point: pH = pK_a (for weak acid). Buffer region exists where pH ≈ pK_a ± 1.

Hydrolysis

Anions of weak acids create basic solutions. Cations of weak bases create acidic solutions. Salts with both act like buffers if Ka ≈ Kb for the relevant species.

Polyprotic Acids

Each proton has its own K_a. K_a1 >> K_a2 >> K_a3. Most calculations only consider the first dissociation unless specifically asked otherwise.

Electrochemistry: Electrons on the Move

Redox Basics

Oxidation = loss of electrons. Reduction = gain of electrons. The oxidizing agent gets reduced. The reducing agent gets oxidized.

Use oxidation numbers to track electrons. They change in redox reactions.

Galvanic vs. Electrolytic Cells

• Galvanic (voltaic): spontaneous, E°_cell > 0, produces electricity
• Electrolytic: nonspontaneous, requires external energy, drives nonspontaneous reactions

Both have two half-cells. Anode is where oxidation happens. Cathode is where reduction happens. Electrons flow from anode to cathode. Cations move toward the cathode. Anions move toward the anode.

Standard Cell Potential

E°_cell = E°_cathode - E°_anode

Use standard reduction potentials. Higher E° = stronger oxidizing agent (more likely to be reduced).

ΔG° = -nFE°

n = moles of electrons transferred. F = 96,485 C/mol. If E° is positive, ΔG° is negative, and the reaction is spontaneous.

ΔG° = -RT ln K

So E° = (RT/nF) ln K or E° = (0.0592/n) log K at 25°C.

Nernst Equation

For nonstandard conditions:

E = E° - (RT/nF) ln Q

Or at 25°C: E = E° - (0.0592/n) log Q

As the reaction proceeds, E decreases. At equilibrium, E = 0 and Q = K.

Electrolysis Applications

Electrolytic cells plate metals, refine copper, and produce chemicals like chlorine. Calculate moles of product using stoichiometry from the balanced half-reaction. 1 Faraday (96,485 C) = 1 mole of electrons.

Laboratory Skills: What the FRQs Actually Test

The AP exam tests lab reasoning directly. You need to interpret data, identify errors, and suggest improvements.

Common Lab Concepts

• Significant figures: Every measurement has limits. Report only what you know for certain plus one estimated digit.
• Uncertainty: Equipment has inherent uncertainty. A buret with ±0.05 mL precision means your volume readings have that margin.
• Error analysis: Systematic errors affect accuracy (off by a consistent amount). Random errors affect precision (scattered results).
• Calorimetry calculations: q_solution = mcΔT. Assume solution has water's specific heat (4.18 J/g·°C) unless told otherwise.
• Titration precision: Run trials until you get concordant results (within 0.2 mL of each other).

Reading Graphs

Be ready to extract data from titration curves, Beer-Lambert plots, kinetic plots, and heating curves. Know how to calculate slope and what it represents.

How to Actually Use This Packet

Reading isn't enough. Here's what you do:

1. Identify gaps. Mark sections where you're unsure. These are your priorities.
2. Practice calculations. Every equation here needs muscle memory. Do 10 problems per equation until it's automatic.
3. Use past FRQs. The College Board releases them. Do them under timed conditions. Compare your answers to scoring guidelines.
4. Know your constants. R = 0.0821, F = 96,485, Kw = 10^-14, 0.0592, 0.693. Write them down when you start the free response.
5. Check units constantly. Wrong units = wrong answer. Convert to moles, liters, Kelvin, and atmospheres before plugging into equations.

What the Exam Actually Expects

The multiple choice tests breadth. You need to recognize concepts and apply them quickly. The free response tests depth. You need to explain reasoning, set up calculations correctly, and justify answers.

For free response:

• Show your work. The graders want to see your thought process.
• Round only at the end. Keep extra sig figs in intermediate steps.
• Answer every part. Even if you bomb one part, the next might be independent.
• Write legibly. Garbled handwriting gets zero credit.

The exam is hard. The pass rate hovers around 50%. But the questions follow patterns. You can prepare for them. Use this packet as a checklist. Cover a section. Test yourself. Repeat until everything sticks.