AP Chemistry Cheat Sheet 2026
Complete concept reference for AP Chemistry — reaction types, acid-base shortcuts, thermodynamics, kinetics, equilibrium, and electrochemistry. Different from the formula sheet: this covers how to use the concepts, not just the equations.
⚗️ Reaction Types
| Type | Pattern | Key Notes |
|---|---|---|
| Synthesis | A + B → AB | Two substances combine. Always exothermic when forming ionic compounds from elements. |
| Decomposition | AB → A + B | One compound breaks apart. Electrolysis, heat (carbonate → CO₂ + oxide), or light. |
| Single Replacement | A + BC → AC + B | Use the activity series. Metal only replaces a metal below it; halogen replaces a less reactive halogen. |
| Double Replacement | AB + CD → AD + CB | Check solubility rules for precipitate. Check acid-base tables for neutralization. |
| Combustion | CₓHᵧ + O₂ → CO₂ + H₂O | Always produces CO₂ and H₂O (complete combustion). CO if O₂ is limited. |
| Acid-Base (Neutralization) | HA + BOH → BA + H₂O | Strong acid + strong base → salt + water. Net ionic: H⁺ + OH⁻ → H₂O |
| Oxidation-Reduction | Transfer of electrons | OIL RIG: Oxidation Is Loss, Reduction Is Gain. Assign oxidation states first. |
🧪 Acid-Base Shortcuts
Strong Acids and Bases (memorize — these fully dissociate)
Strong Acids: HCl, HBr, HI, H₂SO₄, HNO₃, HClO₄, HClO₃
Strong Bases: Group 1 hydroxides (NaOH, KOH, LiOH, etc.) + Ca(OH)₂, Sr(OH)₂, Ba(OH)₂
Strong Bases: Group 1 hydroxides (NaOH, KOH, LiOH, etc.) + Ca(OH)₂, Sr(OH)₂, Ba(OH)₂
| Concept | Formula / Rule | Notes |
|---|---|---|
| pH definition | pH = −log[H⁺] | pH + pOH = 14 at 25°C |
| Strong acid pH | pH = −log[HA] | 100% dissociation; [H⁺] = [HA]₀ |
| Weak acid pH | ICE table → Ka = x²/(C − x) | If Ka is small: x ≈ √(Ka × C) |
| Buffer pH | Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]) | Best buffer when [A⁻] = [HA], so pH = pKa |
| Titration equivalence | moles acid = moles base at equivalence | For weak acid/strong base: pH > 7 at equivalence (conjugate base) |
| Ka × Kb | Ka × Kb = Kw = 1×10⁻¹⁴ | Use to find Kb if Ka of conjugate acid is given (or vice versa) |
| Amphoteric species | Water, HSO₄⁻, HCO₃⁻, HPO₄²⁻, H₂PO₄⁻ | Can act as both acid and base |
⚖️ Equilibrium
| Concept | Key Rule |
|---|---|
| K expression | K = [products]^n / [reactants]^m. Solids and pure liquids are OMITTED (activity = 1). Units: none. |
| Q vs K | Q < K: reaction proceeds forward (makes more product). Q > K: reaction proceeds reverse. Q = K: at equilibrium. |
| Le Chatelier's Principle | Adding reactant → shifts right. Removing product → shifts right. Increasing pressure → shifts toward fewer moles of gas. Increasing temperature → shifts toward endothermic direction. |
| Effect of catalyst | Lowers activation energy — speeds up BOTH directions equally. Does NOT shift equilibrium; Keq unchanged. |
| Kp vs Kc | Kp = Kc(RT)^Δn, where Δn = moles gas products − moles gas reactants |
| Ksp (solubility product) | For MₓAᵧ: Ksp = [M^(y+)]^x [A^(x-)]^y. Compound precipitates when Q > Ksp. |
| Common ion effect | Adding a common ion decreases solubility (shifts equilibrium backward, suppresses dissociation). |
🌡️ Thermodynamics
| Quantity | Sign Rules & Meaning |
|---|---|
| ΔH (enthalpy) | Negative = exothermic (releases heat). Positive = endothermic (absorbs heat). ΔH°rxn = Σ ΔHf°(products) − Σ ΔHf°(reactants) |
| ΔS (entropy) | Positive = more disorder. Gas > liquid > solid. More moles of gas = more entropy. Dissolving ionic salts usually increases entropy. |
| ΔG (Gibbs free energy) | ΔG = ΔH − TΔS. Negative ΔG = spontaneous. ΔG°rxn = Σ ΔGf°(products) − Σ ΔGf°(reactants) |
| Spontaneity summary | ΔH(−) + ΔS(+): always spontaneous. ΔH(+) + ΔS(−): never spontaneous. ΔH(−) + ΔS(−): spontaneous at low T. ΔH(+) + ΔS(+): spontaneous at high T. |
| ΔG and Keq | ΔG° = −RT ln K. If K > 1: ΔG° < 0 (products favored). If K < 1: ΔG° > 0 (reactants favored). |
| Hess's Law | ΔH of a reaction = sum of ΔH of steps. Reversing a step flips the sign of ΔH. Multiplying by a coefficient multiplies ΔH. |
⚡ Kinetics
| Concept | Key Rule |
|---|---|
| Rate law | Rate = k[A]^m[B]^n. Order (m, n) determined ONLY by experiment — not from stoichiometry. |
| Reaction order from data | Compare two experiments where one concentration is held constant. If doubling [A] doubles the rate → 1st order in A. Quadruples → 2nd order. No change → 0th order. |
| Integrated rate laws | 0th: [A] = [A]₀ − kt. 1st: ln[A] = ln[A]₀ − kt. 2nd: 1/[A] = 1/[A]₀ + kt |
| Half-life shortcuts | 1st order: t½ = 0.693/k (constant, independent of concentration). 0th order: t½ = [A]₀/2k. 2nd order: t½ = 1/(k[A]₀) |
| Arrhenius equation | k = Ae^(−Ea/RT). Higher Ea = slower reaction. Higher T = faster reaction. Catalyst lowers Ea without changing ΔH. |
| Collision theory | Molecules must collide with sufficient energy (≥ Ea) AND correct orientation. Increasing T raises fraction of molecules with sufficient energy. |
| Rate-determining step | Slowest step in a mechanism. The rate law for the overall reaction matches the rate law of this step. |
🔋 Electrochemistry
| Concept | Key Rule |
|---|---|
| Galvanic cell | Spontaneous redox reaction (ΔG < 0). Anode = oxidation (loses electrons). Cathode = reduction (gains electrons). Electrons flow: anode → cathode through wire. |
| Cell potential | E°cell = E°cathode − E°anode. Positive E°cell = spontaneous. |
| ΔG and E°cell | ΔG° = −nFE°, where n = moles electrons transferred, F = 96,485 C/mol. |
| Nernst equation | E = E° − (RT/nF) ln Q. At 25°C: E = E° − (0.0592/n) log Q |
| Electrolytic cell | Non-spontaneous redox driven by external power supply. Still: anode = oxidation, cathode = reduction. Used for electroplating and electrolysis of water. |
| Activity series link | Metals higher in the activity series have more negative standard reduction potentials (harder to reduce = easier to oxidize). |
Memory trick: "AN OX, RED CAT" — ANode = OXidation, REDuction at CAThode. Works for BOTH galvanic AND electrolytic cells.
🔬 Atomic Structure & Periodic Trends
| Trend | Direction | Reason |
|---|---|---|
| Atomic radius | Increases down group, decreases across period (left to right) | More protons → stronger nuclear pull on same-shell electrons |
| Ionization energy | Decreases down group, increases across period | Larger radius = electrons farther from nucleus = easier to remove |
| Electron affinity | Generally increases across period (more negative = more favorable) | More protons = stronger attraction for additional electron; halogens highest |
| Electronegativity | Same as ionization energy (F = highest) | F is the most electronegative element (3.98 on Pauling scale) |
| Metallic character | Increases down group, decreases across period | More electron shells + lower ionization energy = easier to lose electrons |
🔗 Bonding & IMFs Quick Reference
| Bond / IMF Type | Relative Strength | When It Occurs |
|---|---|---|
| Ionic bond | Strongest | Metal + nonmetal; large electronegativity difference (>1.7) |
| Covalent bond (polar) | Strong | Nonmetals with electronegativity difference 0.4–1.7 |
| Covalent bond (nonpolar) | Strong | Same elements or difference < 0.4 |
| Hydrogen bond | Moderate | H bonded to N, O, or F. Explains water's anomalously high bp. |
| Dipole-dipole | Weak | Polar molecules (no H–N/O/F required) |
| London dispersion | Weakest | ALL molecules. Increases with molar mass and surface area. |
Boiling point rule: Higher IMF strength → higher boiling point. Same IMF type → larger molecule has higher bp (more LDF). H-bonding explains why H₂O (bp 100°C) boils much higher than H₂S (bp −60°C).