AP Chemistry Practice Test 20
60 original exam-style multiple-choice questions with answers & explanations — covering all nine units of the AP Chemistry course.
Note: These are original practice questions written in the AP Chemistry exam style — not reproduced College Board questions. Use them to test your readiness for the real AP Chemistry exam.
📱 How to use: Try each question first, then click “Show Answer” to check. Covers all 9 AP units: atomic structure, compound structure, intermolecular forces, reactions, kinetics, thermodynamics, equilibrium, acids & bases, and applications of thermodynamics.
⚛ Unit 1: Atomic Structure & Properties
Q1Which of the following is the correct electron configuration for a neutral chromium atom?
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Correct: (B)
Chromium is an exception to the expected filling order. Having a half-filled 3d subshell is especially stable, so one 4s electron moves into the 3d subshell, giving [Ar] 4s^1 3d^5.
Q2A photoelectron spectrum of an element shows five peaks, corresponding to subshells containing 2, 2, 6, 2, and 4 electrons. The element is
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Correct: (D)
The subshells contain a total of 2 + 2 + 6 + 2 + 4 = 16 electrons, so the atomic number is 16, which is sulfur (1s^2 2s^2 2p^6 3s^2 3p^4).
Q3Which of the following elements has the largest first ionization energy?
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Correct: (C)
First ionization energy generally increases across a period. Argon, the noble gas at the far right of Period 3, holds its electrons most tightly and has the largest first ionization energy of the choices.
Q4Boron has two naturally occurring isotopes: B-10 with mass 10.01 amu and 19.9% abundance, and B-11 with mass 11.01 amu and 80.1% abundance. The average atomic mass of boron is closest to
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Correct: (A)
(0.199)(10.01) + (0.801)(11.01) = 1.99 + 8.82 = 10.81 amu.
Q5The ions Na^+, Mg^2+, F^-, and O^2- are isoelectronic (each has 10 electrons). Which of these ions is the largest?
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Correct: (B)
For isoelectronic species, size decreases as the number of protons increases, because a larger nuclear charge pulls the same number of electrons inward more strongly. O^2- has the fewest protons (8), so it is the largest.
Q6A mass spectrum of neon shows three isotopes with masses and percent abundances of 20 amu (90.5%), 21 amu (0.3%), and 22 amu (9.2%). The average atomic mass of neon is closest to
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Correct: (D)
(0.905)(20) + (0.003)(21) + (0.092)(22) = 18.10 + 0.06 + 2.02 = 20.19 amu.
Q7An electron travels at a speed of 2.0 × 10^6 m/s. What is its de Broglie wavelength? (h = 6.626 × 10^-34 J·s, mass of an electron = 9.11 × 10^-31 kg)
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Correct: (A)
λ = h/(mv) = (6.626 × 10^-34)/[(9.11 × 10^-31)(2.0 × 10^6)] = 3.6 × 10^-10 m.
🔗 Unit 2: Molecular & Ionic Compound Structure
Q8What is the formal charge on the nitrogen atom in the ammonium ion, NH4^+?
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Correct: (C)
Formal charge = 5 valence electrons − (0 nonbonding electrons + half of 8 bonding electrons) = 5 − 4 = +1.
Q9What is the molecular geometry of ClF3?
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Correct: (D)
ClF3 has five electron domains (three bonding pairs and two lone pairs) in a trigonal bipyramidal arrangement. The two lone pairs occupy equatorial positions, leaving the three F atoms in a T-shaped molecular geometry.
Q10What is the hybridization of the sulfur atom in SF6?
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Correct: (B)
Sulfur in SF6 is surrounded by six bonding pairs in an octahedral geometry, which requires six hybrid orbitals formed from one s, three p, and two d orbitals: sp^3d^2.
Q11Which of the following compounds contains ionic bonds?
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Correct: (A)
CsF is composed of Cs^+ and F^- ions held together by electrostatic attractions. The other substances consist of covalently bonded molecules.
Q12Which of the following compounds has the smallest lattice energy?
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Correct: (C)
Lattice energy decreases as the ions get larger and their charges decrease. RbI has large singly charged ions that are far apart, giving the weakest attractions and the smallest lattice energy.
Q13Which of the following molecules is polar?
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Correct: (B)
NH3 has a trigonal pyramidal shape with a lone pair on nitrogen, so its bond dipoles do not cancel and the molecule is polar. BF3, CCl4, and CO2 are symmetric, so their dipoles cancel.
Q14How many sigma bonds are present in a molecule of ethene, C2H4?
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Correct: (D)
Ethene has four C-H sigma bonds and one C-C sigma bond, totaling 5 sigma bonds. The C=C double bond additionally contains one pi bond.
💧 Unit 3: Intermolecular Forces & Properties
Q15Which of the following substances exhibits hydrogen bonding in its pure liquid form?
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Correct: (A)
CH3NH2 contains N-H bonds, and the nitrogen lone pair allows neighboring molecules to hydrogen bond. CH3OCH3 has no O-H or N-H bond, and CH4 and CH3Cl cannot hydrogen bond.
Q16Which noble gas has the highest boiling point?
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Correct: (C)
Boiling points of noble gases increase with atomic size because larger electron clouds are more polarizable, producing stronger London dispersion forces. Kr is the largest of the listed gases.
Q17How much heat is required to vaporize 34.0 g of ethanol at its boiling point? (ΔHvap of ethanol = 38.6 kJ/mol; molar mass = 46.07 g/mol)
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Correct: (B)
Moles of ethanol = 34.0/46.07 = 0.738 mol. Heat required = (0.738 mol)(38.6 kJ/mol) = 28.5 kJ.
Q18Which pair of liquids is miscible in all proportions?
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Correct: (D)
Ethanol is polar and forms hydrogen bonds with water, so the two liquids mix completely. The nonpolar liquids (hexane, CCl4, oil) do not mix with water because they cannot form favorable interactions with it.
Q19Which of the following solids conducts electricity in the solid state?
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Correct: (A)
In graphite, each carbon atom contributes a delocalized pi electron that can move within the sheet, allowing electrical conduction. In diamond all electrons are localized in bonds, and in solid NaCl the ions are fixed in the lattice.
Q20Which change will increase the solubility of CO2 gas in water?
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Correct: (C)
By Henry's law, the solubility of a gas is directly proportional to its partial pressure above the liquid. Increasing the CO2 pressure forces more gas into solution; heating would decrease gas solubility.
⚖ Unit 4: Chemical Reactions
Q21When the equation Fe2O3 + CO → Fe + CO2 is balanced with smallest whole-number coefficients, the coefficient of CO is
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Correct: (D)
The balanced equation is Fe2O3 + 3 CO → 2 Fe + 3 CO2, so the coefficient of CO is 3.
Q22Solutions of Pb(NO3)2 and KI are mixed and a yellow precipitate forms. The net ionic equation for the reaction is
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Correct: (B)
PbI2 is the insoluble product. K^+ and NO3^- are spectator ions, so the net ionic equation shows only Pb^2+ and I^- forming solid PbI2.
Q23Potassium chlorate decomposes according to the equation 2 KClO3 → 2 KCl + 3 O2. What mass of O2 is produced when 24.5 g of KClO3 decomposes completely? (Molar mass of KClO3 = 122.55 g/mol, O2 = 32.00 g/mol)
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Correct: (A)
24.5 g KClO3 is 0.200 mol. The mole ratio gives 0.200 × (3/2) = 0.300 mol O2, which has a mass of 0.300 × 32.00 = 9.60 g.
Q24Ammonia is produced by the reaction N2 + 3 H2 → 2 NH3. If 1.00 mol of N2 is mixed with 2.50 mol of H2, how many moles of NH3 can be produced?
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Correct: (D)
The 1.00 mol of N2 would require 3.00 mol of H2, but only 2.50 mol is available, so H2 is the limiting reactant. NH3 produced = 2.50 × (2/3) = 1.67 mol.
Q25A student carries out a reaction with a theoretical yield of 25.0 g and obtains 19.5 g of purified product. The percent yield is
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Correct: (C)
Percent yield = (19.5/25.0) × 100% = 78.0%.
Q26A 20.0 mL sample of H2SO4 solution is titrated with 0.100 M NaOH, and 30.0 mL of the NaOH is required to reach the equivalence point. What is the concentration of the H2SO4? (H2SO4 reacts with 2 NaOH)
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Correct: (B)
Moles of NaOH = (0.0300 L)(0.100 M) = 3.00 × 10^-3 mol. Moles of H2SO4 = half of this = 1.50 × 10^-3 mol. Concentration = 1.50 × 10^-3/0.0200 L = 0.0750 M.
Q27What is the oxidation number of manganese in KMnO4?
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Correct: (A)
With potassium at +1 and each oxygen at −2, the sum of oxidation numbers must be zero: +1 + Mn + 4(−2) = 0, so Mn = +7.
⏱ Unit 5: Kinetics
Q28Initial-rate data for the reaction A + B → products are shown below.
Experiment 1: [A] = 0.20 M, [B] = 0.10 M, rate = 0.010 M/s
Experiment 2: [A] = 0.20 M, [B] = 0.20 M, rate = 0.040 M/s
Experiment 3: [A] = 0.40 M, [B] = 0.10 M, rate = 0.020 M/s
What is the rate law for the reaction?
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Correct: (D)
Doubling [B] (Expt 1 to 2) quadruples the rate, so the reaction is second order in B. Doubling [A] (Expt 1 to 3) doubles the rate, so it is first order in A. Rate = k[A][B]^2.
Q29A radioactive sample undergoes first-order decay. What fraction of the original sample remains after three half-lives?
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Correct: (C)
Each half-life reduces the amount by half: (1/2)^3 = 1/8 = 12.5% of the original sample remains.
Q30What are the units of the rate constant k for a reaction that is second order overall?
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Correct: (B)
Rate has units of M/s. For Rate = k[A]^2, k must have units of (M/s)/M^2 = M^-1·s^-1.
Q31The decomposition of hydrogen peroxide in the presence of iodide follows the mechanism:
Step 1 (slow): H2O2 + I^- → H2O + IO^-
Step 2 (fast): H2O2 + IO^- → H2O + O2 + I^-
Which species is an intermediate in this mechanism?
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Correct: (D)
IO^- is produced in step 1 and consumed in step 2, which makes it an intermediate. I^- is a catalyst because it is consumed in step 1 and regenerated in step 2.
Q32A plot of ln k versus 1/T for a reaction is a straight line with a slope of −6000 K. What is the activation energy of the reaction? (R = 8.314 J/mol·K)
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Correct: (A)
The slope of an Arrhenius plot equals −Ea/R. Therefore Ea = (6000 K)(8.314 J/mol·K) = 4.99 × 10^4 J/mol = 49.9 kJ/mol.
Q33Increasing the concentration of a reactant usually increases the rate of a reaction because
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Correct: (C)
A higher concentration means more molecules are present in the same volume, so collisions between reactant molecules happen more often and the reaction proceeds faster.
Q34On a reaction energy diagram, the reactants are at 50 kJ/mol, the transition state is at 120 kJ/mol, and the products are at 20 kJ/mol. What is the activation energy of the reverse reaction?
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Correct: (B)
The reverse activation energy is measured from the products up to the transition state: 120 − 20 = 100 kJ/mol. (The forward Ea is 70 kJ/mol and ΔH = −30 kJ/mol.)
🔥 Unit 6: Thermodynamics
Q35A reaction is carried out in a calorimeter containing 200.0 g of water, and the temperature of the water rises by 5.0°C. How much heat did the water absorb? (Specific heat of water = 4.18 J/g·°C)
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Correct: (D)
q = mcΔT = (200.0 g)(4.18 J/g·°C)(5.0°C) = 4180 J = 4.18 kJ.
Q36Given the following thermochemical equations:
H2(g) + 1/2 O2(g) → H2O(l), ΔH = −286 kJ
H2O2(l) → H2O(l) + 1/2 O2(g), ΔH = −98 kJ
What is ΔH for the reaction H2(g) + O2(g) → H2O2(l)?
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Correct: (A)
Reverse the second equation (ΔH = +98 kJ: H2O + 1/2 O2 → H2O2) and add it to the first equation. The H2O cancels: ΔH = −286 + 98 = −188 kJ.
Q37Use average bond enthalpies to estimate ΔH for the combustion of methane: CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g). (C-H = 413 kJ/mol, O=O = 498 kJ/mol, C=O = 799 kJ/mol, O-H = 463 kJ/mol)
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Correct: (C)
Bonds broken: 4(413) + 2(498) = 1652 + 996 = 2648 kJ. Bonds formed: 2(799) + 4(463) = 1598 + 1852 = 3450 kJ. ΔH = 2648 − 3450 = −802 kJ.
Q38Using standard enthalpies of formation (ΔHf°: SO2(g) = −296.8 kJ/mol, SO3(g) = −395.7 kJ/mol), the enthalpy change for the reaction 2 SO2(g) + O2(g) → 2 SO3(g) is
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Correct: (A)
ΔH = 2(−395.7) − [2(−296.8) + 0] = −791.4 + 593.6 = −197.8 kJ.
Q39Which gas has the greatest standard molar entropy at 298 K? (Approximate S° values may be inferred from molecular structure.)
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Correct: (B)
Standard molar entropy increases with molecular complexity because larger molecules have more atoms, more bonds, and more available microstates. C2H6 (S° ≈ 230 J/mol·K) exceeds CO2 (≈ 214), CO (≈ 198), and CH4 (≈ 186).
Q40A 25.0 g sample of a substance absorbs 500 J of heat and its temperature increases by 10.0°C. What is the specific heat capacity of the substance?
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Correct: (D)
c = q/(mΔT) = 500/(25.0 × 10.0) = 500/250 = 2.00 J/g·°C.
Q41In a coffee-cup calorimeter, 50.0 mL of 1.0 M HCl is mixed with 50.0 mL of 1.0 M NaOH. The total solution mass is 100.0 g and the temperature rises by 6.9°C. What is the molar enthalpy of neutralization? (Specific heat of the solution = 4.18 J/g·°C)
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Correct: (C)
Heat released = (100.0 g)(4.18)(6.9) = 2884 J, produced by 0.0500 mol of reaction. ΔH = −2884 J/0.0500 mol = −5.77 × 10^4 J/mol = −57.7 kJ/mol. The value is negative because the reaction is exothermic.
⇆ Unit 7: Equilibrium
Q42For the equilibrium CaCO3(s) ⇌ CaO(s) + CO2(g), the correct expression for Kc is
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Correct: (B)
Pure solids are omitted from equilibrium constant expressions because their concentrations do not change. Only the gas CO2 appears, so Kc = [CO2].
Q43The reaction N2O4(g) ⇌ 2 NO2(g) has Kc = 0.200 at a certain temperature. If a flask initially contains 0.500 M N2O4 and no NO2, what is the equilibrium concentration of NO2?
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Correct: (A)
From an ICE table, (2x)^2/(0.500 − x) = 0.200, which gives 4x^2 + 0.200x − 0.100 = 0 and x = 0.135. Therefore [NO2] = 2x = 0.270 M.
Q44The equilibrium PCl5(g) ⇌ PCl3(g) + Cl2(g) is endothermic in the forward direction. If the temperature of the equilibrium mixture is increased, the equilibrium will
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Correct: (D)
For an endothermic reaction, heat can be treated as a reactant. Adding heat shifts the equilibrium toward products, and because more product forms, the value of Kc increases.
Q45The Ksp of Ca(OH)2 is 5.5 × 10^-6. What is the molar solubility of Ca(OH)2 in pure water?
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Correct: (C)
Ca(OH)2 dissolves to give [Ca^2+] = s and [OH^-] = 2s, so Ksp = (s)(2s)^2 = 4s^3 = 5.5 × 10^-6. Then s^3 = 1.4 × 10^-6 and s = 1.1 × 10^-2 M.
Q46A solution is prepared with [Ag^+] = 1.0 × 10^-6 M and [Cl^-] = 1.0 × 10^-4 M. Given that Ksp of AgCl = 1.8 × 10^-10, which of the following will occur?
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Correct: (B)
The ion product Q = (1.0 × 10^-6)(1.0 × 10^-4) = 1.0 × 10^-10, which is less than Ksp = 1.8 × 10^-10. Since Q < Ksp, no precipitate forms.
Q47What is the molar solubility of Mg(OH)2 in pure water? (Ksp = 1.8 × 10^-11)
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Correct: (A)
Ksp = (s)(2s)^2 = 4s^3 = 1.8 × 10^-11, so s^3 = 4.5 × 10^-12 and s = 1.7 × 10^-4 M.
Q48At equilibrium for the reaction 2 NO2(g) ⇌ N2O4(g), the concentrations are [NO2] = 0.020 M and [N2O4] = 0.040 M. What is the value of Kc?
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Correct: (D)
Kc = [N2O4]/[NO2]^2 = 0.040/(0.020)^2 = 0.040/(4.0 × 10^-4) = 100.
🧫 Unit 8: Acids & Bases
Q49What is the pH of a 0.050 M NaOH solution?
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Correct: (C)
NaOH is a strong base, so [OH^-] = 0.050 M. pOH = −log(0.050) = 1.30, and pH = 14.00 − 1.30 = 12.70.
Q50What is the pH of a 0.10 M solution of NH3? (Kb of NH3 = 1.8 × 10^-5)
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Correct: (A)
[OH^-] ≈ (Kb × C)^1/2 = (1.8 × 10^-5 × 0.10)^1/2 = 1.3 × 10^-3 M. Then pOH = 2.87 and pH = 14.00 − 2.87 = 11.13.
Q51A buffer is made by mixing equal concentrations of NH4Cl and NH3. If the pKa of NH4^+ is 9.25, the pH of the buffer is
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Correct: (B)
By the Henderson-Hasselbalch equation, pH = pKa + log([base]/[acid]). With equal concentrations, the log term is log(1) = 0, so pH = pKa = 9.25.
Q52What is the conjugate acid of the carbonate ion, CO3^2-?
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Correct: (D)
A conjugate acid is formed by adding one proton to a base. Adding H^+ to CO3^2- gives HCO3^-.
Q53Acetic acid (pKa = 4.74) is titrated with NaOH. At the half-equivalence point of the titration, the pH of the solution is
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Correct: (C)
At the half-equivalence point, half of the acid has been converted to its conjugate base, so [HA] = [A^-]. The Henderson-Hasselbalch equation then gives pH = pKa = 4.74.
Q54Based on their Ka values, which of the following is the strongest acid?
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Correct: (A)
Acid strength increases with Ka. HF has the largest Ka (6.8 × 10^-4), slightly greater than HNO2 (4.5 × 10^-4) and much greater than CH3COOH and HCN.
🔋 Unit 9: Applications of Thermodynamics
Q55A process has ΔH = +50 kJ and ΔS = +100 J/K. At what temperature does the process change from nonspontaneous to spontaneous (i.e., ΔG = 0)?
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Correct: (B)
Setting ΔG = ΔH − TΔS = 0 gives T = ΔH/ΔS = 50,000 J/(100 J/K) = 500 K. Above 500 K, the TΔS term dominates and the process is spontaneous.
Q56A galvanic cell is constructed from a Ag^+/Ag half-cell (E° = +0.80 V) and a Fe^2+/Fe half-cell (E° = −0.44 V). What is the standard cell potential?
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Correct: (D)
Silver(I) is reduced at the cathode (+0.80 V) and iron is oxidized at the anode (−0.44 V). E°cell = 0.80 − (−0.44) = +1.24 V.
Q57A galvanic cell reaction that transfers 2 mol of electrons has ΔG° = −193 kJ. What is the standard cell potential? (F = 96,485 C/mol e^-)
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Correct: (A)
E° = −ΔG°/(nF) = 193,000/[(2)(96,485)] = +1.00 V. The positive sign is consistent with a spontaneous (galvanic) reaction.
Q58A current of 1.00 A is passed through a solution of AgNO3 for 30.0 min (1800 s). What mass of silver is deposited at the cathode? (F = 96,485 C/mol e^-; molar mass of Ag = 107.87 g/mol)
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Correct: (C)
Charge passed = It = (1.00)(1800) = 1800 C, which is 1800/96,485 = 0.0187 mol of electrons. Since Ag^+ + e^- → Ag, 0.0187 mol of Ag forms, with mass 0.0187 × 107.87 = 2.01 g.
Q59For which combination of signs of ΔH and ΔS will a process be spontaneous at all temperatures?
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Correct: (B)
With ΔH negative and ΔS positive, both terms in ΔG = ΔH − TΔS are negative, so ΔG < 0 at every temperature.
Q60Which of the following statements about an electrolytic cell is correct?
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Correct: (C)
An electrolytic cell uses electrical energy from an external source to force a reaction with ΔG > 0 to occur. As in all cells, reduction occurs at the cathode and oxidation at the anode, and ions must be present to carry charge.
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