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AP Chemistry Practice Exam 2

based on 17 ratings
By — McGraw-Hill Professional
Updated on Feb 9, 2011

Below is a practice exam for AP chemistry.  There are two sections in this practice exam.  Section I has 75 multiple choice questions.  Section II has 6 free response questions.  For a thorough review of the concepts in this practice example, refer to the information center on AP Chemistry Notes.

Multiple Choice

Time—1 hour and 30 minutes

Answer the following questions in the time allowed. You may use the periodic table at the back of the book.

  1. Choose the strongest Lewis acid from the following.
    1. SO42–
    2. Fe3+
    3. Cl
    4. Na+
    5. NO3
  2. Chlorine forms a number of oxyacids. Which of the following is the correct order of increasing acid strength?
    1. HClO3 < HClO4 < HClO < HClO2
    2. HClO4 < HClO3 < HClO2 < HClO
    3. HClO < HClO2 < HClO3 < HClO4
    4. HClO4 < HClO3 = HClO2 < HClO
    5. HClO4 = HClO3 < HClO2 < HClO
  3. Oxidation of which of the following substances will yield a stronger acid?
    1. H2CO3
    2. HNO3
    3. HIO
    4. HIO4
    5. H2SO4
  4. A student prepared five vinegar samples by pipetting 10.00-mL samples of vinegar into five separate beakers. Each of the samples was diluted with deionized water, and phenolphthalein was added as an indicator. The samples were then titrated with standard sodium hydroxide until the appearance of a permanent pink color indicated the end point of the titration. The following volumes were obtained.
  5. Volumes of standard NaOH

    Sample 1: 43.28 mL

    Sample 2: 43.27 mL

    Sample 3: 50.00 mL no color change

    Sample 4: 43.26 mL

    Sample 5: 43.24 mL

    Which of the following is the most likely cause for the variation in the results?

    1. Too much deionized water was added to the third sample.
    2. The student forgot to add phenolphthalein to the third sample.
    3. More acetic acid was present in the third vinegar sample.
    4. The student did not properly rinse the buret with sodium hydroxide solution.
    5. Too much indicator was in the third sample.
  6. Choose the group that does not contain isotopes of the same element.
  7. Choose the ion with the largest ionic radius.
    1. S2–
    2. H+
    3. Li+
    4. O2–
    5. Mg2+
  8. You are given an aqueous solution of aCl2. The simplest method for the separation of BaCl2 from the solution is
    1. filtration of the solution
    2. osmosis of the solution
    3. evaporation of the solution to dryness
    4. electrolysis of the solution
    5. centrifuging the solution
  9. Choose one of the following for questions 8–10.

    1. Cu2+
    2. CO32–
    3. Fe3+
    4. Al3+
    5. Pb2+
  10. This ion is amphoteric.
  11. This ion gives a white precipitate when added to a calcium nitrate solution.
  12. Solutions containing this ion give a reddishbrown precipitate upon standing.
  13. Which of the following is the correct net ionic equation for the addition of aqueous potassium sulfate to a solution of barium chloride?
    1. Ba2+ + SO4 2– → BaSO4
    2. Ba2+ + K2SO4 → BaSO4 + 2 K+
    3. BaCl2 + K2SO4 → BaSO4 + 2 KCl
    4. BaCl2 + SO4 2– → BaSO4 + 2 Cl
    5. BaCl2 + P2SO4 → BaSO4+ 2 PCl
  14. A solution is prepared for qualitative analysis. The solution contains the following ions: Pb2+, Ni2+, and Al3+. Which of the following will cause no observable reaction?
    1. Dilute HCl(aq) is added.
    2. Dilute K2CrO4(aq) is added.
    3. Dilute KCl(aq) is added.
    4. Dilute NH3(aq) is added.
    5. Dilute HNO3(aq) is added.
  15. The addition of concentrated BaCl2(aq) to a 1.0 M (NH4)2SO4 solution will result in which of the following observations?
    1. Nothing happens because the two solutions are immiscible.
    2. The formation of a white precipitate takes place.
    3. An odorless gas forms and bubbles out of the mixture.
    4. The solution becomes basic.
    5. The odor of ammonia will be detected.
  16. How many milliliters of 0.300 M H2SO4 are required to neutralize 50.0 mL of 0.600 M KOH?
    1. 20.0 mL
    2. 25.0 mL
    3. 30.0 mL
    4. 60.0 mL
    5. 50.0 mL
  17. How many grams of hydrogen are in 25 g of (NH4)2SO4?
    1. 2.0 g
    2. 8.0 g
    3. 1.5 g
    4. 3.0 g
    5. 0.75 g
  18. Silver oxide, Ag2O, can be decomposed to silver metal, Ag, plus oxygen gas, O2. How many moles of oxygen gas will form when 4.64 g of solid silver oxide is decomposed? The formula mass of silver oxide is 232.
    1. 0.100 mol
    2. 0.0100 mol
    3. 0.0200 mol
    4. 0.0150 mol
    5. 0.0250 mol
  19. _____ C4H11N(l) + ____ O2(g) → ____ CO2(g) + ____ H2O(l)+ ____ N2(g)
  20. When the above equation is balanced, the lowest whole-number coefficient for N2 is:

    1. 2
    2. 22
    3. 16
    4. 27
    5. 4
  21. 2 KMnO4 + 5 H2C2O4 + 3 H2SO4 → K2SO4 + 2 MnSO4 + 10 CO2 + 8 H2O
  22. How many moles of MnSO4 are produced when 2.0 mol of KMnO4, 10 mol of H2C2O4, and 6.0 mol of H2SO4 are mixed?

    1. 1.0 mol
    2. 2.0 mol
    3. 3.5 mol
    4. 2.5 mol
    5. 3.0 mol
  23. When the following equation is balanced, it is found that 2.00 mol of C8H18 reacts with how many moles of O2?
  24. ____ C8H18+ ____ O2 → ____ CO2 + ____ H2O

    1. 37.5 mol
    2. 2.00 mol
    3. 25.0 mol
    4. 1.00 mol
    5. 12.5 mol
  25. Ra + 2 H2O → Ra(OH)2 + H2
  26. Radium reacts with water according to the above reaction. What volume of hydrogen gas, at standard temperature and pressure, is produced from 0.0100 mol of radium?

    1. 0.560 L
    2. 0.224 L
    3. 0.448 L
    4. 0.112 L
    5. 0.336 L
  27. Three flexible containers are used for gases. The containers are at the same temperature and pressure. One container has 2.0 g of hydrogen, another has 32.0 g of oxygen and the third has 44.0 g of carbon dioxide. Pick the false statement from the following list:
    1. The densities increase in the order hydrogen < oxygen < carbon dioxide.
    2. The number of molecules in all the containers is the same.
    3. The average kinetic energy of all the molecules is the same.
    4. The volume of all three containers is the same.
    5. The average speed of all the molecules is the same.
  28. If a sample of SO2 effuses at a rate of 0.0035 mol per hour at 20°C, which of the gases below will effuse at approximately double the rate under the same conditions?
    1. H2
    2. O2
    3. CO
    4. He
    5. CH4
  29. Choose from the following types of energy for questions 23– 26.

    1. free energy
    2. lattice energy
    3. kinetic energy
    4. activation energy
    5. ionization energy
  30. The energy required to produce a gaseous cation from a gaseous atom in the ground state
  31. The average ______________ is the same for any ideal gas at a given temperature.
  32. The maximum energy available for useful work from a spontaneous reaction
  33. The energy required to separate cations from anions in an ionic solid
  34. When cerium(III) acetate, Ce(C2H3O2)3, is dissolved in water, the temperature increases. Which of the following conclusions may be related to this?
    1. The hydration energies of cerium(III) ions and acetate ions are very low.
    2. Cerium(III) acetate is less soluble in hot water.
    3. The solution is not an ideal solution.
    4. The heat of solution for cerium(III) acetate is endothermic.
    5. The lattice energy of cerium(III) acetate is very low.
  35. Choose the reaction expected to have the greatest increase in entropy.
    1. 2 H2(g) + O2(g) → 2 H2O(g)
    2. 2 Mn2O7(l) → 4 MnO2(s) + 3 O2(g)
    3. C(s) + O2(g) → CO2(g)
    4. 2 Ca(s) + O2(g) → 2 CaO(s)
    5. NH3(g) → NH3(l)
  36. A certain reaction is nonspontaneous under standard conditions, but becomes spontaneous at lower temperatures. What conclusions may be drawn under standard conditions?
    1. ΔH < 0, ΔS < 0 and ΔG = 0
    2. ΔH > 0, ΔS < 0 and ΔG > 0
    3. ΔH < 0, ΔS > 0 and ΔG > 0
    4. ΔH > 0, ΔS > 0 and ΔG > 0
    5. ΔH < 0, ΔS < 0 and ΔG > 0
  37. 4 NO2(g) + Oz(g) → 2 N2O5(g) ΔH = –111 kJ
  38. Determine ΔH for the above reaction if N2O5(s) were formed in the above reaction instead of N2O5(g). The ΔH of sublimation for N2O5 is 54 kJ/mol.

    1. +54 kJ
    2. +219 kJ
    3. +165 kJ
    4. –219 kJ
    5. –165 kJ
  39. Which of the following groups contains only atoms that are diamagnetic in their ground state?
    1. He, Co, and Sr
    2. Zn, Mg, and Xe
    3. O, Be, and Ne
    4. Ca, Mn, and Ar
    5. As, Ba, and Rn
  40. The following ground-state electron configurations are to be used for questions 32–35:

    1. 1s21p62s22p3
    2. 1s22s22p63s23p64s23dl04p65s24d1
    3. 1s22s22p63s23p63d3
    4. 1s22s22p5
    5. 1s22s22p63s23p64s23d104p6
  41. It is not possible for this electron configuration to exist.
  42. A halogen has this electron configuration.
  43. A transition metal atom might have this configuration.
  44. A transition metal ion could have this configuration.
  45. The following answers are to be used for questions 36–39:

    1. Pauli exclusion principle
    2. electron shielding
    3. the wave properties of matter
    4. Heisenberg uncertainty principle
    5. Hund's rule
  46. The exact position of an electron is not known.
  47. Nitrogen atoms, in their ground state, are paramagnetic.
  48. An atomic orbital can hold no more than two electrons.
  49. The 4s orbital fills before the 3d.
  50. Which of the following does not have one or more π bonds?
    1. HNO2
    2. N2
    3. N2H4
    4. HNO3
    5. N2H2
  51. For questions 41 and 42, pick the best choice from the following:

    1. ionic bonds
    2. hybrid orbitals
    3. resonance structures
    4. hydrogen bonding
    5. van der Waals attractions
  52. The unusually high melting point of hydrogen fluoride is due to this.
  53. Their presence explains why the bonds in BF3 are all the same.
  54. Which of the following has more than one unshared pair of valence electrons on the central atom?
  55. Which types of hybridization of carbon are in the compound propane, CH3CH2CH3?
    1. sp3
    2. sp2
    3. sp
    1. I only
    2. II and III
    3. I, II, and III
    4. II only
    5. I and II
  56. The approximate boiling points for hydrogen compounds of some of the elements in the nitrogen family are: (SbH3 15°C), (AsH3 –62°C), (PH3 –87°C), and (NH3, –33°C). The best explanation for the fact that NH3 does not follow the trend of the other hydrogen compounds is
    1. NH3 is the only one that is nearly ideal in the gas phase.
    2. NH3 is the only one that is a base.
    3. NH3 is the only one that is water-soluble.
    4. NH3 is the only one to exhibit hydrogen bonding.
    5. NH3 is the only one that is nonpolar.
  57. Choose the appropriate answer from the following list for questions 46 and 47.

    1. London dispersion forces
    2. covalent bonding
    3. hydrogen bonding
    4. metallic bonding
    5. ionic bonding
  58. This is why copper is ductile.
  59. This is why acetic acid molecules exist as dimmers in the gaseous phase.
  60. Which of the following best explains why 1-butanol, CH3CH2CH2CH2OH, has a higher boiling point (117°C) than its isomer, methyl propyl ether, CH3OCH2CH2CH3 (39°C)?
    1. the lack of hydrogen bonding in 1-butanol
    2. the presence of hydrogen bonding in 1-butanol
    3. the higher molecular mass of 1-butanol
    4. the lower specific heat of 1-butanol
    5. the higher density of 1-butanol
  61. What additional information is needed to convert the molality of a 1.00 m Na2SO4 solution to the molarity?
    1. the volume of the solution
    2. the density of the solution
    3. the boiling point of the solution
    4. the osmotic pressure of the solution
    5. the mass of the solution
  62. All the following substances will dissolve in water. Pick the nonelectrolyte.
    1. Ca(NO3)2
    2. KOH
    3. C2H5OH
    4. HCl
    5. CH3COOH
  63. To prepare 4.0 L of a 0.50-molar KClO3 solution (molecular mass 122.6), a student should follow which of the following procedures?
    1. The student should weigh 245.2 g of solute and add 4.0 L of water.
    2. The student should weigh 61.3 g of solute and add sufficient water to obtain a final volume of 4.0 L.
    3. The student should weigh 61.3 g of solute and add 4.0 Kg of water.
    4. The student should weigh 245.2 g of solute and add sufficient water to obtain a final volume of 4.0 L.
    5. The student should weigh 61.3 g of solute and add 4.0 L of water.
  64. Choose the aqueous solution with the highest boiling point.
    1. 0.20 M HNO3
    2. 0.20 M HClO
    3. 0.40 M C3H7OH
    4. 0.20 M KI
    5. 0.20 M Na2SO4
  65. Which of the following aqueous solutions freezes at the lowest temperature?
    1. 0.25 m NaNO3
    2. 0.25 m FeSO4
    3. 0.25 m C12H22O11
    4. 0.25 m (NH4)2CrO4
    5. 0.25 m KCl
  66. For the following reaction, H2(g) + I2(g) → 2 HI(g), the rate law is: Rate = k[H2][I2]. If a small amount of iodine vapor (I2) is added to a reaction mixture that was 0.10 molar in H2 and 0.20 molar in I2, which of the following statements is true?
    1. Both k and the reaction rate decrease.
    2. Both k and the reaction rate increase.
    3. Both k and the reaction rate remain the same.
    4. Only k increases; the reaction rate remains the same.
    5. Only the rate increases; k remains the same.
  67. Step 1: (CH3)3CBr(aq) → (CH3)3C+(aq) + Br(aq)
  68. Step 2: (CH3)3C+(aq) + H2O(l) → CH3)3 COH2 +(aq)

    Step 3: (CH3)3COH2 +(aq) → H+(aq) + (CH3)3 COH(aq)

    The above represents a proposed mechanism for the hydrolysis of (CH3)3CBr. What are the overall products of the reaction?

    1. (CH3)3C+and Br–
    2. (CH3)3COH2 +and H+
    3. (CH3)3COH and H+
    4. (CH3)3COH and H+
    5. H+ and Br
  69. The table below gives the initial concentrations and rate for three experiments.
  70. The reaction is H2(g) + 2 NO(g) → N2O(g) + H2O(g). What is the rate law for this reaction?

    1. Rate = k[NO]
    2. Rate = k[NO]2[H2]2
    3. Rate = k[H2]
    4. Rate = k[NO]2[H2]
    5. Rate = k[NO][H2]
  71. A solution of a weak base is titrated with a solution of a standard strong acid. The progress of the titration is followed with a pH meter. Which of the following observations would occur?
    1. At the equivalence point, the pH is below 7.
    2. The pH of the solution gradually decreases throughout the experiment.
    3. At the equivalence point, the pH is 7.
    4. The pOH at the equivalence point equals the pKb of the base.
    5. After the equivalence point, the pH becomes constant because this is the buffer region.
  72. When potassium carbonate is dissolved in water
    1. the solution is neutral.
    2. the solution is basic because of hydrolysis of the CO32– ion.
    3. the solution is basic because of hydrolysis of the K+ ion.
    4. the solution is acidic because of hydrolysis of the K+ion.
    5. the solution is acidic because of hydrolysis of the CO32– ion.
  73. Determine the OH(aq) concentration in 0.10 M pyridine (C5H5N) solution. (Kb for pyridine is 9 × 10–9.)
    1. 9 × 10–9 M
    2. 5 × 10–6 M
    3. 7 × 10–3 M
    4. 1 × 10–1 M
    5. 3 × 10–5 M
  74. FeS(s) + 2 H+ (aq) Fe2+ (aq) + H2S(aq)
  75. What is the equilibrium constant for the above reaction? The successive acid dissociation constants for H2S are 9.5 × 10–8 (Ka1) and 1 × 10–19 (Ka2). Ksp, the solubility product constant, for FeS equals 5.0 × 10–18.

    1. 9.5 × 10–27/5.0 × 10–18
    2. 5.0 × 10–18/ 9.5 × 10–27
    3. 5.0 × 10–18/9.5 × 10–8
    4. 9.5 × 10–8/5.0 × 10–18
    5. 1 × 10–19/5.0 × 10–18
  76. C2H2 (g) + H2O(g) CH3CHO(g) exothermic
  77. An equilibrium mixture of the reactants is placed in a sealed container at 150°C. The amount of the product may be increased by which of the following changes?

    1. adding 1 mol of Ar(g) to the container
    2. decreasing the volume of the container
    3. raising the temperature of the container
    1. I only
    2. II and III
    3. III only
    4. I and II
    5. II only
  78. Ksp for Cr(OH)3 is 1.6 × 10–30. What is the molar solubility of this compound in water?
  79. Choose one of the following for questions 63–66.

    1. There is no change in the voltage.
    2. The voltage becomes zero.
    3. The voltage increases.
    4. The voltage decreases, but stays positive.
    5. The voltage becomes negative.

    The following reaction takes place in a voltaic cell:

    Zn(s) + Cu2+(1 M) → Cu(s) + Zn2+(1 M)

    The cell has its voltage measured and found to be +1.10 volts.

  80. What happens to the voltage when deionized water is added to the zinc compartment?
  81. What happens to the cell voltage when the copper electrode is made larger?
  82. What happens to the cell voltage when the salt bridge is replaced with a zinc wire?
  83. What happens to the cell voltage after the cell has operated for 15 min?
  84. How many moles of Cr may be deposited on the cathode when 0.60 F of electricity is passed through a 1.0 M solution of Cr3+?
    1. 0.60 mol
    2. 0.80 mol
    3. 1.0 mol
    4. 0.30 mol
    5. 0.20 mol
  85. Given the above standard reduction potentials, estimate the approximate value of the equilibrium constant for the following reaction:

    1. 104
    2. 1016
    3. 10–16
    4. 10–8
    5. 108
  86. The reducing agent in the above reaction is which of the following?

    1. Bi3+
    2. SnO22–
    3. Bi
    4. H2O
    5. OH

    Questions 70 and 71 are concerned with the following half-reaction in an electrolytic cell:

      2 IO3 + 6H2O + 10e → I2 + 12 OH
  87. Choose the correct statement from the following list.
    1. This reaction occurs at the anode.
    2. The iodine is reduced from +5 to –1.
    3. The iodine is oxidized from –1 to 0.
    4. This reaction occurs at the cathode.
    5. Water is a catalyst.
  88. If a current of 7.50 amp is passed through the electrolytic cell for 0.45 h, how should you calculate the grams of I2 to form?
    1. (7.50)(0.45)(3600)(253.8)/(10)
    2. (7.50)(0.45)(3600)(126.9)/ (96500)(10)
    3. (7.50)(0.45)(60)(253.8)/(96500)(10)
    4. (7.50)(0.45)(3600)(253.8)/ (96500)(10)
    5. (7.50)(0.45)(126.9)/(96500)(10)
  89. The formation of Th from U occurs by
    1. electron capture
    2. α decay
    3. β decay
    4. position decay
    5. γ decay
  90. What is the missing product in the following nuclear reaction?
  91. The above compound would be classified as

    1. an aldehyde
    2. a ketone
    3. an ester
    4. a carboxylic acid
    5. an alcohol
  92. The transition state is higher in energy than the reactants by an amount called
    1. the heat of reaction
    2. the reaction energy
    3. the free energy
    4. the kinetic energy
    5. the activation energy

STOP. End of AP Chemistry Practice Exam 2—Multiple Choice

Time—1 hour and 35 minutes

Answer the following questions in the time allowed. You may use the tables at the back of the book. Write the answers on a separate sheet of paper.

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