AP Chemistry Practice Test 14

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

Q1Chromium is an exception to the expected orbital filling order. Its ground-state electron configuration is

(A) [Ar] 4s^2 3d^4
(B) [Ar] 4s^1 3d^5
(C) [Ar] 3d^6
(D) [Ar] 4s^2 3d^5
Show Answer

Correct: (B)

A half-filled 3d subshell is especially stable, so one 4s electron is used to complete it, giving [Ar] 4s^1 3d^5.

Q2A photoelectron spectrum of an element shows five peaks with relative areas 2, 2, 6, 2, and 4. The element is

(A) S
(B) P
(C) Cl
(D) Si
Show Answer

Correct: (A)

The areas correspond to 1s^2 2s^2 2p^6 3s^2 3p^4, a total of 16 electrons, which is sulfur.

Q3Which of the following elements has the largest atomic radius?

(A) Ca
(B) Ga
(C) K
(D) Ge
Show Answer

Correct: (C)

Atomic radius decreases across a period and increases down a group; K, in group 1 of period 4, is the largest of these elements.

Q4Magnesium consists of three isotopes: Mg-24 (23.985 amu, 78.99%), Mg-25 (24.986 amu, 10.00%), and Mg-26 (25.983 amu, 11.01%). The average atomic mass of magnesium is closest to

(A) 24.99 amu
(B) 25.00 amu
(C) 23.99 amu
(D) 24.31 amu
Show Answer

Correct: (D)

(23.985)(0.7899) + (24.986)(0.1000) + (25.983)(0.1101) = 18.95 + 2.50 + 2.86 = 24.31 amu.

Q5How many moles are contained in 1.81 × 10^24 atoms of iron?

(A) 1.00 mol
(B) 2.00 mol
(C) 3.00 mol
(D) 6.00 mol
Show Answer

Correct: (C)

1.81 × 10^24 / 6.02 × 10^23 = 3.00 mol.

Q6Which element has the greatest second ionization energy?

(A) Mg
(B) Al
(C) K
(D) Na
Show Answer

Correct: (D)

Removing a second electron from Na+ means removing an electron from a stable noble-gas (neon) core, which requires an exceptionally large amount of energy.

Q7What is the electron configuration of S^2-?

(A) [Ne] 3s^2 3p^4
(B) [Ne] 3s^2 3p^6
(C) [Ne] 3s^2 3p^5
(D) [Ar] 4s^2 3d^10 4p^4
Show Answer

Correct: (B)

Sulfur gains two electrons to complete its 3p subshell, giving the electron configuration of argon.

🔗 Unit 2: Molecular & Ionic Compound Structure

Q8What is the formal charge on the nitrogen atom in NH4+?

(A) 0
(B) −1
(C) +1
(D) +2
Show Answer

Correct: (C)

Formal charge = 5 − (0 nonbonding electrons + 4 bonding electrons) = +1.

Q9What is the molecular geometry of XeF4?

(A) Seesaw
(B) Tetrahedral
(C) Octahedral
(D) Square planar
Show Answer

Correct: (D)

XeF4 has six electron domains with two lone pairs positioned opposite each other, giving a square planar molecular geometry.

Q10What is the hybridization of the carbon atom in CO2?

(A) sp
(B) sp^2
(C) sp^3
(D) sp^3d
Show Answer

Correct: (A)

Carbon in CO2 has two electron domains (two double bonds), which requires sp hybridization.

Q11Which of the following compounds has the smallest lattice energy (smallest magnitude)?

(A) MgO
(B) CaO
(C) NaF
(D) KF
Show Answer

Correct: (D)

Lattice energy decreases with larger ions and smaller charges; KF has singly charged ions and the largest ion sizes of these choices.

Q12Which of the following bonds is the shortest?

(A) C≡C in ethyne
(B) C−C in ethane
(C) C=C in ethene
(D) C−O in methanol
Show Answer

Correct: (A)

Bond length decreases as bond order increases, so the triple bond is shorter than the double and single bonds listed.

Q13Which of the following molecules is nonpolar even though it contains polar bonds?

(A) NH3
(B) H2O
(C) BF3
(D) CH3Cl
Show Answer

Correct: (C)

BF3 is trigonal planar and symmetric, so the three B−F bond dipoles cancel each other.

Q14Which of the following bonds has the most ionic character?

(A) H−Cl
(B) Na−Cl
(C) C−H
(D) Cl−Cl
Show Answer

Correct: (B)

Ionic character increases with the difference in electronegativity, which is greatest between Na and Cl.

💧 Unit 3: Intermolecular Forces & Properties

Q15Which of the following substances has the lowest boiling point?

(A) NH3
(B) CH4
(C) H2O
(D) Ne
Show Answer

Correct: (D)

Neon atoms are held in the liquid only by extremely weak dispersion forces, giving it the lowest boiling point (−246°C).

Q16Which of the following substances is most soluble in hexane, a nonpolar solvent?

(A) NaCl
(B) CH3OH
(C) I2
(D) NH3
Show Answer

Correct: (C)

By "like dissolves like," the nonpolar solute I2 dissolves best in the nonpolar solvent hexane.

Q17How much energy is required to melt 36.0 g of ice at 0°C? (ΔHfus = 6.01 kJ/mol)

(A) 12.0 kJ
(B) 3.01 kJ
(C) 6.01 kJ
(D) 24.0 kJ
Show Answer

Correct: (A)

36.0 g / 18.0 g/mol = 2.00 mol; 2.00 mol × 6.01 kJ/mol = 12.0 kJ.

Q18At a given temperature, which of the following liquids has the highest vapor pressure?

(A) Water
(B) Ethanol
(C) Diethyl ether
(D) Ethylene glycol
Show Answer

Correct: (C)

Diethyl ether has the weakest intermolecular forces of these liquids, so its molecules escape into the vapor phase most easily.

Q19Which of the following is a metallic solid?

(A) Dry ice (CO2)
(B) Diamond
(C) NaCl
(D) Copper
Show Answer

Correct: (D)

Copper consists of metal cations surrounded by a sea of delocalized electrons, the defining structure of a metallic solid.

Q20What is the molality of a solution prepared by dissolving 0.250 mol of glucose in 500.0 g of water?

(A) 1.00 m
(B) 0.250 m
(C) 2.00 m
(D) 0.500 m
Show Answer

Correct: (D)

Molality = 0.250 mol / 0.5000 kg of solvent = 0.500 m.

⚖ Unit 4: Chemical Reactions

Q21When the equation NH3(g) + O2(g) → NO(g) + H2O(g) is balanced with the smallest whole-number coefficients, the coefficient of H2O is

(A) 4
(B) 5
(C) 6
(D) 7
Show Answer

Correct: (C)

The balanced equation is 4 NH3 + 5 O2 → 4 NO + 6 H2O.

Q22Solutions of AgNO3(aq) and HCl(aq) are mixed and a white precipitate forms. The net ionic equation for the reaction is

(A) Ag+(aq) + Cl-(aq) → AgCl(s)
(B) Ag+(aq) + NO3-(aq) + H+(aq) + Cl-(aq) → AgCl(s) + H+(aq) + NO3-(aq)
(C) AgNO3(aq) + HCl(aq) → AgCl(s) + HNO3(aq)
(D) H+(aq) + NO3-(aq) → HNO3(aq)
Show Answer

Correct: (A)

AgCl precipitates; H+ and NO3- are spectator ions and are omitted.

Q23What is the oxidation number of chromium in K2Cr2O7?

(A) +2
(B) +3
(C) +7
(D) +6
Show Answer

Correct: (D)

2(+1) + 2(Cr) + 7(−2) = 0 gives 2(Cr) = +12, so Cr = +6.

Q24What mass of H2O is produced by the complete combustion of 16.0 g of CH4? (CH4 + 2 O2 → CO2 + 2 H2O; molar mass of CH4 = 16.0 g/mol)

(A) 18.0 g
(B) 9.00 g
(C) 54.0 g
(D) 36.0 g
Show Answer

Correct: (D)

16.0 g / 16.0 g/mol = 1.00 mol CH4, which produces 2.00 mol H2O; 2.00 × 18.0 = 36.0 g.

Q25Aluminum reacts with chlorine according to 2 Al + 3 Cl2 → 2 AlCl3. If 5.40 g of Al (0.200 mol) reacts with 10.65 g of Cl2 (0.150 mol), what mass of AlCl3 (molar mass 133.3 g/mol) can form?

(A) 13.3 g
(B) 26.7 g
(C) 6.67 g
(D) 40.0 g
Show Answer

Correct: (A)

Cl2 is the limiting reactant: 0.150 mol Cl2 forms (2/3)(0.150) = 0.100 mol AlCl3; 0.100 × 133.3 = 13.3 g.

Q26A reaction produces 18.0 g of product, which corresponds to a percent yield of 75.0%. The theoretical yield of the reaction was

(A) 13.5 g
(B) 18.0 g
(C) 24.0 g
(D) 31.5 g
Show Answer

Correct: (C)

Theoretical yield = actual yield / percent yield = 18.0 g / 0.750 = 24.0 g.

Q27What volume of 0.400 M NaOH is required to neutralize 20.00 mL of 0.200 M H2SO4?

(A) 10.0 mL
(B) 20.0 mL
(C) 40.0 mL
(D) 80.0 mL
Show Answer

Correct: (B)

Moles of H2SO4 = 0.02000 L × 0.200 M = 0.00400 mol; because H2SO4 is diprotic, 0.00800 mol of NaOH is needed; 0.00800 / 0.400 = 0.0200 L = 20.0 mL.

⏱ Unit 5: Kinetics

Q28Rate data for the reaction A + B → products are: Experiment 1: [A] = 0.20 M, [B] = 0.10 M, rate = 3.0 × 10^-3 M/s; Experiment 2: [A] = 0.20 M, [B] = 0.20 M, rate = 1.2 × 10^-2 M/s; Experiment 3: [A] = 0.40 M, [B] = 0.10 M, rate = 6.0 × 10^-3 M/s. The rate law for the reaction is

(A) Rate = k[A][B]
(B) Rate = k[A]^2[B]
(C) Rate = k[A][B]^2
(D) Rate = k[A]^2[B]^2
Show Answer

Correct: (C)

Doubling [B] quadruples the rate (second order in B), while doubling [A] only doubles the rate (first order in A).

Q29The units of the rate constant for a first-order reaction are

(A) M/s
(B) M^-1 s^-1
(C) M^-2 s^-1
(D) s^-1
Show Answer

Correct: (D)

For a first-order reaction, rate = k[A], so k must have units of s^-1.

Q30A first-order reaction has a half-life of 20.0 min. If the initial concentration of reactant is 0.640 M, what is the concentration after 60.0 min?

(A) 0.320 M
(B) 0.160 M
(C) 0.080 M
(D) 0.040 M
Show Answer

Correct: (C)

60.0 min is three half-lives, so the concentration decreases by a factor of 2^3 = 8: 0.640 / 8 = 0.080 M.

Q31A zero-order reaction has k = 0.020 M/s. If the initial concentration of reactant is 0.50 M, how long will it take for the reactant to be completely consumed?

(A) 25 s
(B) 10 s
(C) 50 s
(D) 100 s
Show Answer

Correct: (A)

For zero-order kinetics the concentration decreases at a constant rate: t = 0.50 M / 0.020 M/s = 25 s.

Q32Consider the two-step mechanism: Step 1: A + B → C + D; Step 2: C + E → F + D. Which species is a reaction intermediate?

(A) C
(B) A
(C) D
(D) F
Show Answer

Correct: (A)

C is produced in Step 1 and consumed in Step 2, which is the definition of an intermediate.

Q33According to collision theory, increasing the concentration of reactants increases the reaction rate because

(A) The activation energy decreases
(B) Reactant particles collide more frequently
(C) The particles move more slowly
(D) The enthalpy change of the reaction decreases
Show Answer

Correct: (B)

More particles per unit volume produce more frequent collisions, so more collisions per unit time can overcome the activation energy.

Q34For a certain reaction, the activation energy of the forward reaction is 80 kJ/mol and ΔH = −30 kJ/mol. The activation energy of the reverse reaction is

(A) 50 kJ/mol
(B) 80 kJ/mol
(C) 110 kJ/mol
(D) 30 kJ/mol
Show Answer

Correct: (C)

Ea(reverse) = Ea(forward) − ΔH = 80 − (−30) = 110 kJ/mol.




🔥 Unit 6: Thermodynamics

Q35How much heat is absorbed when 200.0 g of water is heated by 10.0°C? (specific heat of water = 4.18 J/g·°C)

(A) 0.836 kJ
(B) 8.36 kJ
(C) 4.18 kJ
(D) 83.6 kJ
Show Answer

Correct: (B)

q = mcΔT = (200.0 g)(4.18 J/g·°C)(10.0°C) = 8360 J = 8.36 kJ.

Q36Given: S(s) + O2(g) → SO2(g), ΔH = −297 kJ; and 2 SO2(g) + O2(g) → 2 SO3(g), ΔH = −198 kJ. What is ΔH for S(s) + 3/2 O2(g) → SO3(g)?

(A) −99 kJ
(B) −396 kJ
(C) −495 kJ
(D) −297 kJ
Show Answer

Correct: (B)

Add the first equation to one-half of the second: −297 + (−198 / 2) = −297 − 99 = −396 kJ.

Q37Using bond enthalpies (N≡N = 941 kJ/mol, H−H = 436 kJ/mol, N−H = 391 kJ/mol), estimate ΔH for N2(g) + 3 H2(g) → 2 NH3(g).

(A) −97 kJ
(B) +97 kJ
(C) −4595 kJ
(D) −49 kJ
Show Answer

Correct: (A)

Bonds broken: 941 + 3(436) = 2249 kJ; bonds formed: 6(391) = 2346 kJ; ΔH = 2249 − 2346 = −97 kJ.

Q38Given ΔHf°(CO) = −110.5 kJ/mol and ΔHf°(CO2) = −393.5 kJ/mol, what is ΔH for 2 CO(g) + O2(g) → 2 CO2(g)?

(A) −283 kJ
(B) −787 kJ
(C) −1132 kJ
(D) −566 kJ
Show Answer

Correct: (D)

ΔH = 2(−393.5) − 2(−110.5) = −787 + 221 = −566 kJ.

Q39For which of the following reactions is ΔS negative?

(A) N2(g) + 3 H2(g) → 2 NH3(g)
(B) 2 NH3(g) → N2(g) + 3 H2(g)
(C) H2O(l) → H2O(g)
(D) NaCl(s) → Na+(aq) + Cl-(aq)
Show Answer

Correct: (A)

Four moles of gas are converted into two moles of gas, so the disorder and entropy of the system decrease.

Q40A 100.0 g sample of a metal at 100.0°C is placed in 200.0 g of water at 20.0°C, and the final temperature is 24.0°C. What is the specific heat of the metal? (specific heat of water = 4.18 J/g·°C)

(A) 0.220 J/g·°C
(B) 0.440 J/g·°C
(C) 0.880 J/g·°C
(D) 1.76 J/g·°C
Show Answer

Correct: (B)

Heat gained by water = (200.0)(4.18)(4.0) = 3344 J; c(metal) = 3344 / [(100.0)(76.0)] = 0.440 J/g·°C.

Q41A 0.500 mol sample of a fuel is burned in a bomb calorimeter that has a heat capacity of 10.0 kJ/°C, and the temperature of the calorimeter rises by 5.00°C. The enthalpy of combustion of the fuel is

(A) −200 kJ/mol
(B) −50 kJ/mol
(C) −25 kJ/mol
(D) −100 kJ/mol
Show Answer

Correct: (D)

Heat released = (10.0 kJ/°C)(5.00°C) = 50.0 kJ from 0.500 mol, so ΔH = −50.0 / 0.500 = −100 kJ/mol.

⇆ Unit 7: Equilibrium

Q42For the equilibrium CaCO3(s) ⇌ CaO(s) + CO2(g), the equilibrium constant expression Kc is

(A) [CO2]
(B) [CaO][CO2] / [CaCO3]
(C) [CaCO3] / ([CaO][CO2])
(D) [CaO][CO2]
Show Answer

Correct: (A)

Pure solids are omitted from equilibrium constant expressions, leaving Kc = [CO2].

Q43For the reaction 2 NO2(g) ⇌ N2O4(g), Kc = 170 at 298 K. What is Kp at 298 K? (R = 0.0821 L·atm/mol·K)

(A) 6.9
(B) 170
(C) 4.2 × 10^3
(D) 0.14
Show Answer

Correct: (A)

Δn = 1 − 2 = −1, so Kp = Kc(RT)^-1 = 170 / (0.0821 × 298) = 170 / 24.5 = 6.9.

Q44For PCl5(g) ⇌ PCl3(g) + Cl2(g), Kc = 0.040. If a flask initially contains 0.20 M PCl5 and no products, what is [Cl2] at equilibrium?

(A) 0.040 M
(B) 0.072 M
(C) 0.089 M
(D) 0.13 M
Show Answer

Correct: (B)

Setting x = [Cl2] formed, x^2 / (0.20 − x) = 0.040 gives x^2 + 0.040x − 0.0080 = 0, so x = 0.072 M.

Q45For the equilibrium 2 SO3(g) ⇌ 2 SO2(g) + O2(g), if the volume of the container is decreased at constant temperature, the equilibrium will

(A) Not shift because the temperature is constant
(B) Shift right, toward more moles of gas
(C) Shift in a way that increases Kc
(D) Shift left, toward fewer moles of gas
Show Answer

Correct: (D)

Decreasing the volume increases the pressure, so the equilibrium shifts toward the side with fewer moles of gas (2 mol on the left vs. 3 mol on the right).

Q46The molar solubility of PbCl2 in water is 1.6 × 10^-2 M. What is the Ksp of PbCl2?

(A) 1.6 × 10^-5
(B) 4.1 × 10^-6
(C) 2.6 × 10^-4
(D) 6.6 × 10^-5
Show Answer

Correct: (A)

For PbCl2(s) ⇌ Pb^2+(aq) + 2 Cl-(aq), Ksp = (s)(2s)^2 = 4s^3 = 4(1.6 × 10^-2)^3 = 1.6 × 10^-5.

Q47The Ksp of BaSO4 is 1.1 × 10^-10. A solution contains [Ba^2+] = 1.0 × 10^-4 M and [SO4^2-] = 1.0 × 10^-4 M. Which statement is correct?

(A) No precipitate forms because Q is less than Ksp
(B) A precipitate forms because Q is greater than Ksp
(C) The solution is exactly saturated because Q equals Ksp
(D) Additional BaSO4 will dissolve
Show Answer

Correct: (B)

Q = (1.0 × 10^-4)(1.0 × 10^-4) = 1.0 × 10^-8, which is greater than Ksp, so BaSO4 precipitates until Q = Ksp.

Q48The solubility of Mg(OH)2 in 0.10 M MgCl2(aq), compared with its solubility in pure water, is

(A) The same in both solutions
(B) Higher in the MgCl2 solution
(C) Lower in the MgCl2 solution
(D) Zero in both solutions because Mg(OH)2 is insoluble
Show Answer

Correct: (C)

The Mg^2+ common ion shifts Mg(OH)2(s) ⇌ Mg^2+(aq) + 2 OH-(aq) to the left, lowering the solubility.

🧫 Unit 8: Acids & Bases

Q49What is the pH of 0.050 M NaOH(aq)?

(A) 1.30
(B) 12.70
(C) 13.00
(D) 11.70
Show Answer

Correct: (B)

pOH = −log(0.050) = 1.30, so pH = 14.00 − 1.30 = 12.70.

Q50What is the pH of 0.100 M NH3(aq)? (Kb of NH3 = 1.8 × 10^-5)

(A) 11.13
(B) 8.87
(C) 2.87
(D) 10.13
Show Answer

Correct: (A)

[OH-] = sqrt(1.8 × 10^-5 × 0.100) = 1.34 × 10^-3 M; pOH = 2.87, so pH = 14.00 − 2.87 = 11.13.

Q51A buffer solution contains 0.30 M HCOOH and 0.10 M HCOONa. What is the pH of the buffer? (pKa of HCOOH = 3.74)

(A) 3.74
(B) 4.22
(C) 2.74
(D) 3.26
Show Answer

Correct: (D)

pH = pKa + log([base]/[acid]) = 3.74 + log(0.10 / 0.30) = 3.74 − 0.48 = 3.26.

Q52The conjugate acid of HPO4^2- is

(A) H2PO4-
(B) PO4^3-
(C) H3PO4
(D) OH-
Show Answer

Correct: (A)

Adding a proton (H+) to HPO4^2- gives its conjugate acid, H2PO4-.

Q53A weak acid is titrated with a strong base. At the half-equivalence point of the titration,

(A) pH = 7.00
(B) pH = pKb of the conjugate base
(C) pH = 14.00 − pKa
(D) pH = pKa of the weak acid
Show Answer

Correct: (D)

At the half-equivalence point, [acid] = [conjugate base], so the Henderson-Hasselbalch equation gives pH = pKa.

Q54Carbonic acid, H2CO3, has Ka1 = 4.3 × 10^-7 and Ka2 = 5.6 × 10^-11. What is the Kb of HCO3- when it acts as a base (HCO3- + H2O ⇌ H2CO3 + OH-)?

(A) 2.3 × 10^-8
(B) 1.8 × 10^-4
(C) 5.6 × 10^-11
(D) 4.3 × 10^-7
Show Answer

Correct: (A)

This hydrolysis is the reverse of the first ionization, so Kb = Kw / Ka1 = 1.0 × 10^-14 / 4.3 × 10^-7 = 2.3 × 10^-8.

🔋 Unit 9: Applications of Thermodynamics

Q55A reaction is spontaneous at all temperatures when

(A) ΔH > 0 and ΔS > 0
(B) ΔH < 0 and ΔS < 0
(C) ΔH < 0 and ΔS > 0
(D) ΔH > 0 and ΔS < 0
Show Answer

Correct: (C)

With ΔH negative and ΔS positive, ΔG = ΔH − TΔS is negative at every temperature.

Q56A reaction has ΔH° = +50 kJ/mol and ΔS° = +200 J/mol·K. What is ΔG° at 298 K, and is the reaction spontaneous under standard conditions?

(A) +9.6 kJ/mol; nonspontaneous
(B) −9.6 kJ/mol; spontaneous
(C) +109.6 kJ/mol; nonspontaneous
(D) −109.6 kJ/mol; spontaneous
Show Answer

Correct: (B)

ΔG = 50 − (298)(0.200) = 50 − 59.6 = −9.6 kJ/mol; a negative ΔG indicates a spontaneous reaction.

Q57Standard reduction potentials are: Li+/Li = −3.04 V, Zn^2+/Zn = −0.76 V, Cu^2+/Cu = +0.34 V, Ag+/Ag = +0.80 V. Which metal is the strongest reducing agent?

(A) Ag
(B) Cu
(C) Zn
(D) Li
Show Answer

Correct: (D)

The metal with the most negative reduction potential is the most easily oxidized, making Li the strongest reducing agent.

Q58Silver is plated from a solution containing Ag+ using a current of 5.00 A. How long will it take to deposit 10.8 g of silver? (Ag+ + e- → Ag; molar mass of Ag = 107.9 g/mol; F = 96485 C/mol e-)

(A) 965 s
(B) 1930 s
(C) 3860 s
(D) 9649 s
Show Answer

Correct: (B)

10.8 g / 107.9 g/mol = 0.100 mol Ag requires 0.100 mol of electrons = 9650 C; t = 9650 C / 5.00 A = 1930 s.

Q59For a galvanic cell with E°cell = +0.50 V in which 1 mol of electrons is transferred, ΔG° is

(A) −96 kJ/mol
(B) −24 kJ/mol
(C) −48 kJ/mol
(D) +48 kJ/mol
Show Answer

Correct: (C)

ΔG° = −nFE° = −(1)(96485)(0.50) = −48,000 J/mol = −48 kJ/mol.

Q60Which of the following statements is true for any spontaneous process?

(A) ΔH must be negative
(B) The entropy of the universe increases
(C) ΔG is greater than zero
(D) ΔS of the system must be positive
Show Answer

Correct: (B)

By the second law of thermodynamics, every spontaneous process increases the entropy of the universe; ΔH and ΔS of the system alone do not determine spontaneity.