How to review effectively
Before the exam, active recall (saying the answer yourself before checking it) works far better than rereading: wherever you get stuck while recalling is where you’d lose points on the exam.
- Test first, then read: go through the cards below unit by unit and mark anything you can’t answer, or get wrong, as “missed.”
- Split time by weight: Unit 3 (18–22%) and Unit 8 (11–15%) carry the most points; the other units are 7–9% each.
- Go back to the lessons: open each unit with misses, reread that lesson’s examples, and redo its practice.
- Test again a few days later: repeat the same cards after two or three days until you can answer them all.
The core of the nine units
U1 Atomic Structure and Properties 7–9%
In a mass spectrum, what do the horizontal position and relative height of each peak show?
The position is the isotope’s mass (m/z) and the height its relative abundance; average atomic mass = Σ(isotope mass × abundance).
Across a period from left to right, how do atomic radius and first ionization energy change, and why?
The radius decreases and the ionization energy generally increases: the nuclear charge grows while electrons stay in the same shell, so the effective nuclear charge pulls the outer electrons more strongly.
In a photoelectron spectrum (PES), what do a peak’s position and height show?
The position is the binding energy (larger closer to the nucleus) and the height the number of electrons in that subshell.
What are the steps to find an empirical formula from the mass of each element?
Mass → moles → divide each by the smallest → convert to the simplest whole-number ratio.
U2 Compound Structure and Properties 7–9%
How do you use VSEPR to find a molecule’s shape?
Count the electron domains around the central atom (each single, double or triple bond counts as one, and so does each lone pair), find the electron-domain geometry, then name the shape from the atom positions only.
How is formal charge calculated, and how do you pick the best Lewis structure?
FC = valence electrons − nonbonding electrons − bonds; keep formal charges as close to 0 as possible, with any negative formal charge on the more electronegative atom.
How does bond order relate to bond length and bond energy?
The higher the bond order, the shorter and stronger the bond.
What are the melting points and conductivity of ionic, metallic, covalent network and molecular solids?
Ionic solids melt high and conduct when molten or dissolved; metals conduct and are malleable; covalent network solids melt very high and generally don’t conduct; molecular solids melt low and don’t conduct.
U3 Properties of Substances and Mixtures 18–22%
In what order should you consider intermolecular forces when comparing boiling points?
First look for ions or hydrogen bonding, then polarity (dipole–dipole), and finally dispersion forces (more electrons and more surface contact mean stronger dispersion).
State the ideal gas law and Dalton’s law of partial pressures.
PV = nRT; Ptotal = P₁ + P₂ + …, and each partial pressure Pᵢ = χᵢPtotal.
Under what conditions do real gases deviate most from ideal behavior?
At high pressure (the molecules’ own volume matters) and low temperature (attractions between molecules matter).
How is a photon’s energy related to its wavelength, and what is the Beer-Lambert law?
E = hν = hc/λ, so shorter wavelengths carry more energy; A = εbc.
U4 Chemical Reactions 7–9%
What are the steps for writing a net ionic equation?
Split soluble strong electrolytes into ions and cancel spectator ions on both sides; keep precipitates, gases, weak electrolytes and water as formulas.
How do you identify the oxidizing and reducing agents in a redox reaction?
The species whose oxidation number decreases is reduced and is the oxidizing agent; the one whose oxidation number increases is oxidized and is the reducing agent.
What is the basic path for a stoichiometry calculation?
Mass → moles → convert with the equation’s coefficients → moles → mass; with two reactants, find the limiting reactant first.
How is the equivalence point of a titration defined?
The point where the titrant added has reacted exactly and completely with the analyte, according to the mole ratio in the equation.
U5 Kinetics 7–9%
How are the orders in a rate law determined?
Only from experimental data; a rate law can be written from the coefficients only for an elementary step.
What are the integrated rate law and half-life for a first-order reaction?
ln[A] = ln[A]₀ − kt; t½ = 0.693/k, independent of the starting concentration.
How does a catalyst speed up a reaction? Does it change ΔH or K?
It provides a pathway with a lower activation energy; it changes neither ΔH nor K.
In a mechanism, what are the rate-determining step and an intermediate?
The slowest step sets the overall rate law; an intermediate forms in one step and is used up in a later one, so it doesn’t appear in the overall equation.
U6 Thermochemistry 7–9%
What are the three relationships used in calorimetry?
qsoln = mcΔT; qrxn = −qsoln; ΔH = qrxn ÷ n.
How do you find ΔH° from enthalpies of formation?
ΔH° = Σ n·ΔfH°(products) − Σ n·ΔfH°(reactants); elements in their standard states have ΔfH° = 0.
How do you estimate ΔH from bond enthalpies?
ΔH = Σ bond enthalpies of bonds broken − Σ bond enthalpies of bonds formed.
How is Hess’s law used?
The ΔH of the overall reaction is the sum of the steps’ ΔH values; reversing an equation flips the sign of ΔH, and multiplying by a coefficient multiplies ΔH too.
U7 Equilibrium 7–9%
How do you compare Q and K to predict the direction of a reaction?
If Q < K, it proceeds forward; if Q > K, it proceeds in reverse; if Q = K, it’s at equilibrium.
How does K change when the equation changes?
Reversing gives 1/K; multiplying by n gives Kⁿ; adding two equations multiplies their K values.
How is an ICE table used?
Write rows for initial, change and equilibrium concentrations, with changes in the ratio of the coefficients, then substitute into the K expression and solve.
How are Ksp and solubility related, and what is the common-ion effect?
Write the ion concentrations in terms of the solubility s and substitute into Ksp (for AgCl, Ksp = s²); a common ion already in solution lowers the solubility.
U8 Acids and Bases 11–15%
How do you find the pH of a strong acid and of a weak acid?
Strong acid: [H⁺] = c; weak acid: [H⁺] ≈ √(Ka·c) (valid when ionization is under 5%).
How do you find a buffer’s pH, and when is its capacity greatest?
pH = pKa + log([A⁻]/[HA]); capacity is greatest when [A⁻] = [HA], that is, when pH = pKa.
What happens at the half-equivalence and equivalence points of a weak-acid titration curve?
At half-equivalence, pH = pKa; at equivalence the solution contains the conjugate base, so for a weak acid titrated with a strong base the pH is above 7.
How does molecular structure affect acid strength?
The more stable the conjugate base (negative charge on a more electronegative atom, spread out by resonance, or pulled by electron-withdrawing groups), the stronger the acid; a longer, weaker H–X bond also makes a stronger acid.
U9 Thermodynamics and Electrochemistry 7–9%
How do you use ΔG to decide whether a process is thermodynamically favorable?
ΔG = ΔH − TΔS; it’s favorable when ΔG < 0. With ΔH < 0 and ΔS > 0, it’s favorable at every temperature.
How are ΔG° and K related?
ΔG° = −RT ln K; when ΔG° < 0, K > 1.
How do you find E°cell for a galvanic cell, and how is it related to ΔG°?
E°cell = E°cathode − E°anode; ΔG° = −nFE°, so the reaction is favorable when E° > 0.
What is the path for an electrolysis calculation?
q = It → n(e⁻) = q ÷ F → moles of product from the electrons in the half-reaction → mass.
Exam questions often span units
AP questions rarely stay inside one unit. Long FRQs often combine stoichiometry (Unit 4), gases (Unit 3), thermochemistry (Unit 6) and equilibrium (Unit 7) in a single scenario. Each example below links three units.
Example 1: stoichiometry + gases (Units 1, 3, 4)
0.243 g of magnesium reacts with excess hydrochloric acid: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g). What volume of H₂ forms at 25 °C and 0.980 atm? (Mg 24.31 g/mol)
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Moles of magnesium (Unit 1).
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Convert with the equation (Unit 4). n(H₂) : n(Mg) = 1 : 1, so n(H₂) = 1.00 × 10⁻² mol.
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Ideal gas law (Unit 3). With T = 298.15 K:
Example 2: equilibrium + acids + thermodynamics (Units 7, 8, 9)
At 25 °C, HF has Ka = 6.4 × 10⁻⁴. Find ΔG° for the ionization of HF, and the pH of 0.10 M HF.
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ΔG° (Unit 9). ΔG° = −RT ln K:
ΔG° > 0 matches K < 1: only a small fraction of HF ionizes in water.
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Try the approximation first (Unit 8).
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Check it. 8.0 × 10⁻³ ÷ 0.10 = 8%, more than 5%, so the approximation isn’t good enough; solve the quadratic from the ICE table (Unit 7):
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pH. pH = −log(7.7 × 10⁻³) = .
- A catalyst increases a reaction’s equilibrium constant K
A catalyst doesn’t change K; it only speeds up the approach to equilibrium.
- Across a period from left to right, atomic radius decreases
The effective nuclear charge increases.
- When a weak acid is titrated with a strong base, the equivalence pH is 7
At equivalence the solution contains the conjugate base, so the pH is above 7.
- An endothermic reaction can never be favorable at any temperature
If ΔS > 0, ΔG = ΔH − TΔS can be negative at high temperature.
- The higher the bond order, the shorter and stronger the bond
Triple bonds are shorter than double bonds, which are shorter than single bonds.
- A first-order half-life doesn’t depend on the starting concentration
t½ = 0.693/k.
- Lowering the temperature always shifts an equilibrium forward
Only when the forward reaction is exothermic.
- Real gases deviate clearly from ideal behavior at high pressure and low temperature
The molecules’ volume and attractions become significant.
Common mistakes
Mistake 1: “Rereading the book cover to cover counts as review.” Rereading makes everything feel familiar, but the exam asks you to retrieve. Test yourself first, then read what you couldn’t answer.
Mistake 2: “Spend the same time on every unit.” Unit 3 is worth more than twice as much as most units, and Unit 8 clearly more too. Split your time by weight and by your own weak spots.
Mistake 3: “Mastering each unit separately is enough.” Exam questions often span units. In a combined question, you need to recognize which unit’s method each step uses (for example, stoichiometry first, then the gas law).
Quick check
Which unit carries the most weight on the AP Chemistry exam?
Unit 3 is 18–22%, followed by Unit 8 at 11–15%.
Which review method works best?
Active recall exposes weak spots directly.
For a reaction with ΔG° = +18.2 kJ/mol, how does K compare with 1?
ΔG° = −RT ln K, so ΔG° > 0 means ln K < 0.
After approximating x ≈ √(Ka·c), when must you switch to the quadratic?
Above 5% ionization, ignoring x causes too much error.
Which pair of topics is most likely to appear in the same question?
Finding the volume of a gas produced from a reactant’s mass is one of the most common combinations.
Of CH₄, NH₃ and H₂O, which has the highest boiling point?
Each H₂O molecule can form the most hydrogen bonds, so its intermolecular forces are strongest.
AP-style practice
Ammonia synthesis: N₂(g) + 3H₂(g) ⇌ 2NH₃(g). Data at 25 °C (OpenStax Appendix G), with ΔfH° in kJ/mol and S° in J/(mol·K):
| Substance | ΔfH° | S° |
|---|---|---|
| N₂(g) | 0 | 191.6 |
| H₂(g) | 0 | 130.7 |
| NH₃(g) | −45.9 | 192.8 |
(a) Draw the Lewis structure of NH₃, give its shape, and say whether it is polar.
(b) Calculate ΔH° and ΔS° for the reaction.
(c) Calculate ΔG° at 25 °C. Under standard conditions at this temperature, is the forward reaction thermodynamically favored?
(d) Industry runs this reaction at about 700 K. How does raising the temperature affect K and the reaction rate?
Model answer and scoring
(a) N forms a single bond to each of three H atoms and has one lone pair; with four electron domains, the molecule is trigonal pyramidal. The N–H bonds are polar and the pyramidal shape keeps the dipoles from cancelling, so NH₃ is polar.
(b) ΔH° = 2(−45.9) − 0 = −91.8 kJ/mol; ΔS° = 2(192.8) − [191.6 + 3(130.7)] = 385.6 − 583.7 = −198.1 J/(mol·K).
(c) ΔG° = −91.8 − 298.15 × (−0.1981) = −32.7 kJ/mol < 0, so it is thermodynamically favorable under standard conditions at 25 °C.
(d) The forward reaction is exothermic, so raising the temperature makes K smaller (the equilibrium shifts left); but more molecules then have energy above the activation energy, so the rate increases. Industry uses a higher temperature (and a catalyst) so the reaction is fast enough.
Scoring: (a) Lewis structure, trigonal pyramidal, polar; (b) values and signs of ΔH° and ΔS°; (c) ΔG° and the judgment; (d) K decreases and the rate increases, each with a reason.
Summary
- Review with active recall: test yourself, go back to the units you missed, and test again a few days later.
- Split time by weight, with Units 3 and 8 first.
- In combined questions, identify which unit’s method each step uses, and check approximations, units and signs.
AP tip: In the last two weeks, a daily set of questions that span units (such as one long FRQ) is closer to the real exam than reviewing a single unit at a time.
Sources: unit weights from the College Board AP Chemistry Course and Exam Description; Ka of HF from OpenStax Chemistry 2e Appendix H; ΔfH° and S° from Appendix G of the same book.