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AP Chemistry Practice Exam 1

based on 35 ratings
By — McGraw-Hill Professional
Updated on Apr 25, 2014

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 exam, refer to the information center on AP Chemistry Notes.

Multiple Choice

Section 1

Time—1 hour and 30 minutes

NO CALCULATOR MAY BE USED WITH SECTION 1

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

After the above chemical equation is balanced, the lowest whole-number coefficient for water is

Copper metal reacts with nitric acid according to the above equation. A 0.30-mol sample of copper metal and 100.0 mL of 3.0 M nitric acid are mixed in a flask. How many moles of NO gas will form?

When the above equation is balanced, the lowest whole number coefficient for CO2 is:

How many moles of MnSO4 are produced when 2.0 mol of KMnO4, 5.0 mol of H2C2O4, and 1.5 mol of H2SO4 are mixed?

After the above equation is balanced, how many moles of O2 can be produced from 1.0 mol of KClO3?

Strontium reacts with water according to the above reaction. What volume of hydrogen gas, at standard temperature and pressure, is produced from 0.100 mol of strontium?

Choose from the following types of energy for questions 26–28.

ClF(g) + F2(g) → ClF3(l)

Determine ΔH for the above reaction if C2H5OH(l) was formed in the above reaction instead of C2H5OH(g). The ΔH of vaporization for C2H5OH is 43 kJ/mol.

The following answers are to be used for questions 35–38.

Choose from the following solids for questions 45–48.

AP Chemistry Practice Exam 1

Using the above information, choose the best answer for preparing a pH = 7.9 buffer.

Questions 61–64 refer to the following aqueous solutions. All concentrations are 1 M.

Which species, in the above equilibrium, behave as bases?

A 1.00-L flask is filled with 0.70 mol of H2 and 0.60 mol of CO, and allowed to come to equilibrium. At equilibrium, there are 0.40 mol of CO in the flask. What is the value of Kc , the equilibrium constant, for the reaction?

What is the coefficient of OH when the above reaction is balanced?

For the above reaction, pick the true statement from the following:

  1. Choose the strongest Lewis base from the following.
    1. Na+
    2. Fe3+
    3. NH3
    4. Zn2+
    5. BF3
  2. Which of the following CANNOT behave as both a Brønsted base and a Brønsted acid?
    1. HPO42–
    2. C2O42–
    3. HSO4
    4. HC2O4
    5. HCO3
  3. A species, molecule, or ion, is called a Lewis base if it does which of the following?
    1. It is an electron-pair donor.
    2. It donates an H+.
    3. It accepts an H+.
    4. It is an electron-pair acceptor.
    5. It increases the H+(aq) in water.
  4. Which of following are proper laboratory procedures for a titration?
    1. Make sure the color change of the indicator persists for at least 30 s.
    2. Allow all materials to cool to room temperature before they are weighed.
    3. Rinse the buret with deionized water before it is filled with titrant for the first titration.
    1. I and III only
    2. I, II, and III
    3. II only
    4. II and III only
    5. I and II only
  5. In most of its compounds, this element exists as a monatomic cation.
    1. F
    2. S
    3. N
    4. Ca
    5. Cl
  6. In which of the following groups are the species listed correctly in order of increasing radius?
    1. Sr, Ca, Mg
    2. Se2–, S2–, O2–
    3. Mn3+, Mn2+, Mn
    4. I, Br, Cl
    5. K, Ca, Sc
  7. Which of the following elements has the lowest electronegativity?
    1. F
    2. I
    3. K
    4. Al
  8. Which of the following represents the correct formula for hexamminechromium(III) chloride?
    1. [Cr(NH3)6](ClO3)3
    2. (NH3)6Cr3Cl
    3. Am6CrCl3
    4. [Cr(NH3)6]Cl3
    5. [Cr3(NH3)6]Cl3
  9. _____Fe(OH)3(s) + _____H2SeO4(aq) → _____Fe2 (SeO4)3(s) + _____H2O(1)
    1. 1
    2. 6
    3. 9
    4. 12
    5. 3
  10. Which of the following best represents the net ionic equation for the reaction of barium hydroxide with an aqueous potassium sulfate solution?
    1. Ba(OH)2 + KSO4 → BaSO4 + KOH
    2. Ba2+ + K2SO4 → BaSO4 + 2 K+
    3. Ba2+ + SO42– → BaSO4
    4. Ba(OH)2 + SO42– → BaSO4 + 2 OH
    5. Ba(OH)2 + K2SO4 → BaSO4 + 2 KOH
  11. A sample of magnesium metal is heated in the presence of nitrogen gas. After the sample was heated, some water was added to it. Which of the following statements is false?
    1. The magnesium reacted with the nitrogen to produce magnesium nitride.
    2. No reaction occurred because nitrogen gas is so unreactive.
    3. The solid did not dissolve in the water.
    4. After the addition of the water, the distinctive odor of ammonia gas was present.
    5. The water converted some of the magnesium nitride to magnesium hydroxide.
  12. A student mixes 50.0 mL of 0.10 M Ni(NO3)2 solution with 50.0 mL of 0.10 M NaOH. A green precipitate forms, and the concentration of the hydroxide ion becomes very small. Which of the following correctly places the concentrations of the remaining ions in order of decreasing concentration?
    1. [Na+] > [Ni2+] > [NO3]
    2. [Ni2+] > |NO3] > [Na+]
    3. [Na+] > [NO3] > [Ni2+]
    4. [NO3] > [Na+] > [Ni2+]
    5. [Ni2+] > [Na+] > [NO3]
  13. The addition of concentrated Ba(OH)2(aq) to a 1.0 M (NH4)2SO4 solution will result in which of the following observations?
    1. The odor of ammonia is detected, and a white precipitate forms.
    2. The formation of a white precipitate takes place.
    3. The solution becomes acidic.
    4. The odor of ammonia is detected.
    5. An odorless gas forms and bubbles out of the mixture.
  14. Manganese, Mn, forms a number of oxides. A particular oxide is 69.6% Mn. What is the simplest formula for this oxide?
    1. MnO
    2. Mn2O3
    3. Mn3O4
    4. MnO2
    5. Mn2O7
  15. Sodium sulfate forms a number of hydrates. A sample of a hydrate is heated until all the water is removed. What is the formula of the original hydrate if it loses 56% of its mass when heated?
    1. Na2SO4.H2O
    2. Na2SO4.2H2O
    3. Na2SO4.6H2O
    4. Na2SO4.8H2O
    5. Na2SO4.10H2O
  16. 3Cu(s) + 8 HNO3(aq)→ 3 Cu(NO3)2(aq) + 2 NO(g) + 4 H2O(l)
    1. 0.20 mol
    2. 0.038 mol
    3. 0.10 mol
    4. 0.075 mol
    5. 0.30 mol
  17. Gold(III) oxide, Au2O3, can be decomposed to gold metal, Au, plus oxygen gas, O2. How many moles of oxygen gas will form when 2.21 g of solid gold(III) oxide is decomposed? The formula mass of gold(III) oxide is 442.
    1. 0.00750 mol
    2. 0.0150 mol
    3. 0.00500 mol
    4. 0.00250 mol
    5. 0.0100 mol
  18. _____C4H11N(l) + _____O2(g) → _____CO2(g) + _____H2O(l) + _____N2(g)
    1. 4
    2. 16
    3. 27
    4. 22
    5. 2
  19. 2 KMnO4 + 5 H2C2O4 + 3 H2SO4 → K2SO4 + 2 MnSO4 + 10 CO2 + 8 H2O
    1. 2.0 mol
    2. 1.5 mol
    3. 1.0 mol
    4. 3.0 mol
    5. 2.5 mol
  20. _____KClO3 → _____KCl + _____O2
    1. 1.5 mol
    2. 3.0 mol
    3. 1.0 mol
    4. 3.0 mol
    5. 6.0 mol
  21. Sr + 2 H2O → Sr(OH)2 + H2
    1. 3.36 L
    2. 5.60 L
    3. 2.24 L
    4. 4.48 L
    5. 1.12 L
  22. A sample of nitrogen gas is placed in a container with constant volume. The temperature is changed until the pressure doubles. Which of the following also changes?
    1. density
    2. moles
    3. average velocity
    4. number of molecules
    5. potential energy
  23. An experiment to determine the molecular mass of a gas begins by heating a solid to produce a gaseous product. The gas passes through a tube and displaces water in an inverted, water-filled bottle. The mass of the solid is measured, as is the volume and the temperature of the displaced water. Once the barometric pressure has been recorded, what other information is needed to finish the experiment?
    1. the heat of formation of the gas
    2. the density of the water
    3. the mass of the displaced water
    4. the vapor pressure of the water
    5. the temperature to which the solid was heated
  24. Determine the final temperature of a sample of hydrogen gas. The sample initially occupied a volume of 6.00 L at 127°C and 875 mm Hg. The sample was heated, at constant pressure, until it occupied a volume of 15.00 L.
    1. 318°C
    2. 727°C
    3. 45°C
    4. 160°C
    5. 1000°C
  25. From the following, choose the gas that probably shows the least deviation from ideal gas behavior.
    1. Kr
    2. CH4
    3. O2
    4. H2
    5. NH3
    1. free energy
    2. lattice energy
    3. kinetic energy
    4. activation energy
    5. ionization energy
  26. The maximum energy available for useful work from a spontaneous reaction
  27. The energy needed to separate the ions in an ionic solid
  28. The energy difference between the transition state and the reactants
    1. 2 ClF(g) + O2(g) → Cl2O(g) + OF2(g) ΔH° = 167.5 kJ
    2. 2 F2(g) + O2(g) → 2 OF2(g) ΔH° = –43.5 kJ
    3. 2 ClF3(l) + 2 O2(g) → Cl2O(g) + 3 OF2(g) ΔH° = 394.1 kJ
  29. Using the information given above, calculate the enthalpy change for the following reaction:
    1. –135.1 kJ
    2. +135.1 kJ
    3. 270.2 kJ
    4. –270.2 kJ
    5. 0.0 kJ
  30. When lithium sulfate, Li2SO4, is dissolved in water, the temperature increases. Which of the following conclusions may be related to this?
    1. Lithium sulfate is less soluble in hot water.
    2. The hydration energies of lithium ions and sulfate ions are very low.
    3. The heat of solution for lithium sulfate is endothermic.
    4. The solution is not an ideal solution.
    5. The lattice energy of lithium sulfate is very low.
  31. What is the energy required to form a gaseous cation from a gaseous atom?
    1. ionization energy
    2. kinetic energy
    3. activation energy
    4. lattice energy
    5. free energy
  32. C2H4(g) + H2O(g) → C2H5OH(g) ΔH = –46 kJ
    1. +3 kJ
    2. +89 kJ
    3. –3 kJ
    4. +43 kJ
    5. –89 kJ
  33. The ground-state configuration of Ni2+ is which of the following?
    1. 1s22s22p63s23p63d84s2
    2. 1s22s22p63s23p63d104s2
    3. ls22s22p63s23p63d10
    4. ls22s22p63s23p63d8
    5. ls22s22p63s23p63d54s2
  34. A ground-state electron in a calcium atom might have which of the following sets of quantum numbers?
    1. n = 3; l = 2; m1 = 0; ms = –1/2
    2. n = 5; l = 0; m1 = 0; ms = –1/2
    3. n = 4; l = 1; m1 = 0; ms = –1/2
    4. n = 4; l = 0; m1 = 0; ms = –1/2
    5. n = 4; l = 0; m1 = + 1; ms = –1/2
    1. Pauli exclusion principle
    2. electron shielding
    3. the wave properties of matter
    4. Heisenberg uncertainty principle
    5. Hund's rule
  35. The diffraction of electrons
  36. The maximum number of electrons in an atomic orbital is two.
  37. An oxygen atom is paramagnetic in the ground state.
  38. The position and momentum of an electron cannot be determined exactly.
  39. Magnesium reacts with element X to form an ionic compound. If the ground-state electron configuration of X is ls22s22p5, what is the simplest formula for this compound?
    1. Mg2X3
    2. MgX2
    3. MgX4
    4. Mg2X5
    5. MgX
  40. VSEPR predicts that a BF3 molecule will be which of the following shapes?
    1. tetrahedral
    2. trigonal bipyramidal
    3. square pyramid
    4. trigonal planar
    5. square planar
  41. Which of the following is polar?
    1. BF3
    2. IF5
    3. CF4
    4. XeF4
    5. AsF5
  42. The only substance listed below that contains ionic, σ, and π bonds is:
    1. C2H4
    2. NaH
    3. NH4Cl
    4. NaC2H3O2
    5. H2O
  43. Which molecule or ion in the following list has the greatest number of unshared electron pairs around the central atom?
    1. IF7
    2. NO3
    3. BF3
    4. NH3
    5. CBr4
  44. Which of the following processes does not involve breaking an ionic or a covalent bond?
    1. 2 NO(g) + O2 → 2 NO2(g)
    2. NaNO3(s) → Na+(aq) + NO3(aq)
    3. Zn(s) → Zn(g)
    4. 2 H2(g) + O2(g) → 2 H2O(g)
    5. 2 KClO3(s) → 2 KCl(s) + 3 O2(g)
    1. composed of atoms held together by delocalized electrons
    2. composed of molecules held together by intermolecular dipole–dipole interactions
    3. composed of positive and negative ions held together by electrostatic attractions
    4. composed of macromolecules held together by strong bonds
    5. composed of molecules held together by intermolecular London forces
  45. Graphite
  46. Ca(s)
  47. CaCO3(s)
  48. SO2(s)
  49. The critical point represents
    1. the highest temperature and pressure where the substance may exist as discrete solid and gas phases.
    2. the highest temperature and pressure where the substance may exist as discrete liquid and gas phases.
    3. the temperature and pressure where the substance exists in equilibrium as solid, liquid, and gas phases.
    4. the highest temperature and pressure where the substance may exist as discrete liquid and solid phases.
    5. the highest temperature and pressure where a substance can sublime.
  50. A sample of a pure liquid is placed in an open container and heated to the boiling point. Which of the following may increase the boiling point of the liquid?
    1. The container is sealed.
    2. The size of the container is increased.
    3. More liquid is added.
    1. II and III
    2. I and III
    3. III only
    4. II only
    5. I only
  51. Which point on the diagram below might represent the normal boiling point?
  52. What is the total concentration of cations in a solution made by combining 700.0 mL, of 3.0 M (NH4)3PO4 with 300.0 mL of 2.0 M Na2SO4?
    1. 2.7 M
    2. 13 M
    3. 7.5 M
    4. 5.0 M
    5. 2.5 M
  53. A stock solution that is 0.30 M in Na2SO4 is available. How many moles of solid Na3PO4 must be added to 800 mL of this solution to increase the sodium ion concentration to 0.90 M?
    1. 0.060
    2. 0.12
    3. 0.080
    4. 0.16
    5. 0.24
  54. If a solution of ethyl ether, (C2H5)2O, in ethanol, C2H5OH, is treated as an ideal solution, what is the mole fraction of ethyl ether in the vapor over an equimolar solution of these two liquids? The vapor pressure of ethyl ether is 480 mm Hg at 20°C, and the vapor pressure of ethanol is 50 mm Hg at this temperature.
    1. 0.50
    2. 0.76
    3. 0.91
    4. 0.27
    5. 0.09
  55. How many milliliters of concentrated ammonia (7.0-molar NH3) are needed to prepare 0.250 L of 3.0-molar NH3?
    1. 110 mL
    2. 0.11 mL
    3. 200 mL
    4. 150 mL
    5. 75 mL
  56. The plot of ln[A] versus time gives a straight line. This implies the rate law is
    1. rate = k[A]2
    2. rate = k[A]–2
    3. rate = k[A]0
    4. rate = k[A]–1
    5. rate = k[A]
  57. The specific rate constant, k, for radioactive lawrencium-256 is 86 h–1 What mass of a 0.0500 ng sample of lawrencium-256 remains after 58 s?
    1. 0.0500 ng
    2. 0.0250 ng
    3. 0.0125 ng
    4. 0.00625 ng
    5. 0.0375 ng
  58. The purpose of using a lit match to start the fire in a gas grill is
    1. to supply the free energy for the reaction
    2. to catalyze the reaction
    3. to supply the heat of reaction
    4. to supply the kinetic energy for the reaction
    5. to supply the activation energy for the reaction
    1. K2HPO4
    2. K3PO4
    3. K2HPO4 + KH2PO4
    4. K2HPO4 + K3PO4
    5. H3PO4 + KH2PO4
  59. What is the ionization constant, Ka, for a weak monoprotic acid if a 0.6-molar solution has a pH of 2.0?
    1. 1.7 × 10–4
    2. 1.7 × 10–2
    3. 6 × 10–6
    4. 2.7 × 10–3
    5. 3.7 × l0–4
    1. CH3NH2 (methylamine) and LiOH (lithium hydroxide)
    2. C2H5NH2 (ethylamine) and C2H5NH3NO3 (ethylammonium nitrate)
    3. CH3NH2 (methylamine) and HC3H5O2 (propionic acid)
    4. KClO4 (potassium perchlorate) and HClO4 (perchloric acid)
    5. H2C2O4 (oxalic acid) and KHC2O4 (potassium hydrogen oxalate)
  60. The most basic solution (highest pH)
  61. The solution with a pH nearest 7
  62. A buffer with a pH > 7
  63. A buffer with a pH < 7
  64. At constant temperature, a change in volume will NOT affect the moles of the substances present in which of the following?
    1. CO32–
    2. H2O
    3. HCO3
    1. I and III
    2. II only
    3. I and II
    4. I only
    5. II and III
  65. CO(g) + 2H2(g) → CH3OH(g)
    1. 0.74
    2. 3.2
    3. 0.0050
    4. 5.6
    5. 1.2
  66. H2O(l) + CrO42–(aq) + HSnO2(aq) → CrO2 (aq) + OH(aq) + HSnO3(aq)
    1. 10
    2. 2
    3. 5
    4. 4
    5. 1
  67. 2 Bi3+ + 3 SnO22– + 6 OH → 3 SnO32– + 3 H2O + 2 Bi
    1. The oxidation number of tin changes from + 2 to + 4.
    2. The oxidation number of tin changes from + 4 to + 2.
    3. The Bi3+ is oxidized by the tin.
    4. The OH reduces the Bi3+.
    5. The SnO32– is formed by the reduction of SnO22–.
  68. An electrolysis cell was constructed with two platinum electrodes in a 1.00 M aqueous solution of KCl. An odorless gas evolved from one electrode and a gas with a distinctive odor evolved from the other electrode. Choose the correct statement from the following list.
    1. The odorless gas was oxygen.
    2. The odorless gas was evolved at the anode.
    3. The gas with the distinctive odor was evolved at the anode.
    4. The odorless gas was evolved at the positive electrode.
    5. The gas with the distinctive odor was evolved at the negative electrode.
  69. When decays, it emits 2 α particles, then a β particle, followed by an α particle. The resulting nucleus is:
  70. Which of the following lists the types of radiation in the correct order of increasing penetrating power?
    1. α, γ, β
    2. β, α, γ
    3. α, β, γ
    4. β, γ, α
    5. γ, β, α
  71. Which of the following statements are correct concerning β particles?
    1. They have a mass number of zero and a charge of –1.
    2. They are electrons.
    3. They are less penetrating than α particles.
    1. I and II
    2. I and III
    3. II and III
    4. I only
    5. II only
  72. If 75% of a sample of pure decays in 24.6 yr, what is the half-life of ?
    1. 24.6 yr
    2. 18.4 yr
    3. 12.3 yr
    4. 6.15 yr
    5. 3.07 yr
  73. Alkenes are hydrocarbons with the general formula CnH2n. If a 0.420 g sample of any alkene is combusted in excess oxygen, how many moles of water will form?
    1. 0.0600
    2. 0.450
    3. 0.015
    4. 0.300
    5. 0.0300
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