Diagnostic Test

1. Part A

2. Directions

   Each set of lettered choices below refers to the numbered formulas or statements immediately following it. Select the one lettered choice that best fits each formula or statement. In each set, a choice may be used once, more than once, or not at all.

   Note

   For all questions involving solutions and/or chemical equations, you can assume that the system is in water unless otherwise stated.

3. Questions 1–5

4. 1. Atomic mass
   2. Atomic number
   3. Atomic radius
   4. Electronegativity
   5. Ionization potential

5. 1. Takes into account various isotopes of an atom

   2. Determines how electron density is shared when an atom forms a bond

   3. Average distance between the nucleus and the outermost electron

   4. Number of protons in an element

   5. Energy required to remove an electron

6. Questions 6–8

7. 1. Na
   2. Fe
   3. Cl^–
   4. Rb
   5. Ca⁺

8. 6. Has a filled 4p orbital

   7. All of its electrons are paired in this ion

   8. Electron configuration is

      1s²2s²2p⁶3s²3p⁶3d⁶4s²

9. Questions 9–12

10. 1. Cation
    2. Inert Gas
    3. Crystal
    4. Anion
    5. Element

11. 9. Cannot be further broken down by chemical means

    10. Ionic species with a positive charge

    11. Has an octet of valence electrons

    12. Defined by the number of protons

12. Questions 13–16

13. 1. Centrifuge
    2. Barometer
    3. Balance
    4. Calorimeter
    5. Battery

14. 13. Apparatus used to measure the heat absorbed or released by a reaction

    14. Apparatus used to measure atmospheric pressure

    15. Apparatus used to measure weight

    16. Apparatus used to sediment particles in suspension

15. Questions 17–19

16. 1. Gibbs free energy
    2. Heat of formation
    3. Specific heat
    4. Heinsenberg uncertainty principle
    5. Heat of vaporization

17. 17. The energy of a system available to do work

    18. Heat required to raise one unit mass of a substance by 1 degree Celsius

    19. Heat absorbed or released during production of a substance from elements in their standard states

18. Questions 20–23

19. What happens when the following are dissolved in water?

    1. NaOH
    2. CsCl
    3. HBr
    4. CH₃COOH
    5. O₂

20. 20. Forms a strong base

    21. Forms a weak acid

    22. Forms a strong acid

    23. Forms an ionic solution with a neutral pH

1. Part B

2. Directions

   Each question below consists of two statements, statement I in the left-hand column and statement II in the right-hand column. For each question, determine whether statement I is true (T) or false (F) and whether statement II is true (T) or false (F). Then, determine if statement II is a correct explanation (CE) of statement I.

101. I		II
     Diamond and graphite are both substances made up of carbon but have different properties	BECAUSE	they are composed of different isotopes.

102. I		II
     Na and Cl form an ionic bond	BECAUSE	Cl donates an electron to Na.

103. I		II
     AgCl will not dissolve in a concentrated NaCl solution	BECAUSE	the chloride ions from NaCl suppress the solubility of AgCl.

104. I		II
     Hydrogen and deuterium are different elements	BECAUSE	they have a different number of protons.

105. I		II
     On the periodic chart, atomic radius increases from left to right	BECAUSE	the number of protons is increasing.

106. I		II
     An element with an atomic number of X and mass number of N has N–X neutrons	BECAUSE	elements have more neutrons than protons.

107. I		II
     Radioactive decay has a characteristic half-life	BECAUSE	first-order kinetics are found in radioactive decay.

108. I		II
     Sulfur chemically resembles oxygen	BECAUSE	they are in the same period.

109. I		II
     A nonelectrolyte does not ionize in water	BECAUSE	the solution does not conduct electricity.

110. I		II
     SO3 diffuses more slowly than CO2	BECAUSE	it has smaller bond angles.

111. I		II
     Hydrogen bonding is stronger between noble gases than small electronegative atoms such as fluorine	BECAUSE	noble gases have a large dipole moment.

112. I		II
     An electron in a 3s subshell may be excited to jump into the 3p	BECAUSE	the Heisenberg uncertainty principle states that one cannot know what orbital an electron is in.

113. I		II
     A positive ΔG tells you that the reaction is spontaneous	BECAUSE	entropy always decreases in an isolated system.

114. I		II
     An increase in pressure in a closed container with an ideal gas leads to a decrease in volume	BECAUSE	pressure and volume are proportional.

115. I		II
     An amphoteric species acts as either an acid or a base	BECAUSE	it contains both hydrophobic and hydrophilic regions.

116. I		II
     An indicator will allow you to determine whether a solution is acidic or basic	BECAUSE	it will change colors in solutions with different pHs.

Return to the section of your answer sheet you started for chemistry and answer questions 24–69.

1. Part C

2. Directions

   Each of the incomplete statements or questions below is followed by five suggested completions or answers. Select the one that is best for each case.

24. What is the formal charge on the nitrogen atom in HNO₃?

    1. –1
    2. +1
    3. 0
    4. +2
    5. +3

25. What volume of a 1 M solution of hydrochloric acid is required to neutralize 80 mL of a 0.5 M NaOH solution?

    1. 320 mL
    2. 160 mL
    3. 80 mL
    4. 40 mL
    5. 20 mL

26. Which of the following would be classified as a strong electrolyte?

    1. Benzoic acid
    2. Water
    3. Hydrofluoric acid
    4. Potassium chloride
    5. Glucose

27. Which atom has an ionic radius that is larger than its atomic radius?

    1. Na
    2. K
    3. Mg
    4. Ca
    5. Cl

28. The basic structure of crystalline substances is called

    1. unit cell
    2. molecule
    3. lattice
    4. geode
    5. matrix

29. The oxidation state of nitrogen is most negative in which of the following compounds?

    1. N₂
    2. N₂O
    3. NH₃
    4. NO₂
    5. NO₃

30. An insulated tube with a movable piston at one end had 500 J of heat added to it. If, during the experiment, the piston moves and does 75 J of work on the atmosphere, what is the change in energy of the tube system?

    1. 575 J
    2. 500 J
    3. 425 J
    4. –75 J
    5. –425 J

31. The reaction of NaOH and H₂SO₄ in water goes to completion because

    1. it is a neutralization reaction
    2. water is a strong electrolyte
    3. sodium sulfate quickly precipitates
    4. a volatile product is formed
    5. sulfuric acid is a very strong acid

32. Which of the following would be different in a ground state and an excited state neon atom?

    1. The number of neutrons
    2. The number of electrons
    3. The atomic weight
    4. The electronic configuration
    5. Everything would remain the same.

33. The K_sp of Mg(OH)₂ in water is 1.2 × 10^–11. If the Mg²⁺ concentration in an acid solution is 1.2 × 10^–5 mol/L, what is the pH at which Mg(OH)₂ just begins to precipitate?

    1. 3
    2. 4
    3. 5
    4. 11
    5. 12

34. Which of the following states has the highest average translational kinetic energy?

    1. Solid
    2. Liquid
    3. Gas
    4. Colloid
    5. None of the above

35. Which of the following will favor the melting of ice in a closed container if all other parameters are kept constant?

    1. Adding water with a temperature of 0°C
    2. Lowering the temperature below 0°C
    3. Lowering the pressure
    4. Raising the pressure
    5. Decreasing the amount of ice

36. Which of the following has the most polar bond?

    1. N– O
    2. C–H
    3. C– C
    4. H–F
    5. None of the bonds are polar.

37. The rate law expression for the reaction

    N₂ + 3H₂ → 2NH₃
    1. can be represented by rate
    2. can be represented by rate
    3. can be represented by rate
    4. can be represented by rate
    5. cannot be determined from the information given

38. 100 mL of 10 N H₂SO₄ is diluted to 800 mL. What is the molarity of the dilute acid solution?

    1. 16/10 M
    2. 8/10 M
    3. 10/8 M
    4. 10/12 M
    5. 5/8 M

39. Which of the following molecules contains both ionic and covalent bonds?

    1. C₆H₁₄
    2. MgCl₂
    3. (NH₄)₂SO₄
    4. H₂O
    5. C₂H₄

40. Gas A is at 30°C and gas B is at 20°C. Both gases are at 1 atmosphere. What is the ratio of the volume of 1 mole of gas A to 1 mole of gas B?

    1. 1:1
    2. 2:3
    3. 3:2
    4. 303:293
    5. 606:293

41. How will the equilibrium of the following reaction be affected if more chlorine is added?

    PCl₅ (g) ⇌ PCl₃ (g) + Cl₂ (g)
    1. It will be shifted to the right.
    2. It will be shifted to the left.
    3. It will be unaffected.
    4. The effect on the equilibrium cannot be determined without more information.
    5. More PCl₃ will be produced.

42. After balancing the equation

    …BrO₃^– (aq) + …Br^– (aq) + …H⁺(aq) → …Br₂ (l) + …H₂O

    the ratio of BrO₃^– to Br^– is

    1. 1:5
    2. 1:3
    3. 1:2
    4. 1:1
    5. 2:3

43. What mass of sodium carbonate, Na₂CO₃ (formula weight = 106 amu), is needed to make 120 mL of a 1.5 M solution?

    1. 295 g
    2. 9.5 g
    3. 19 g
    4. 589 g
    5. 19,000 g

44. Equimolar amounts of hydrogen and oxygen gas, at the same temperature, are released into a large container. The ratio of the rate of diffusion of the hydrogen molecules to that of the molecules of oxygen would be

    1. 256:1
    2. 16:1
    3. 1:16
    4. 4:1
    5. 1:4

45. What does X represent in the following nuclear reaction?

    
    1. β particle
    2. Positron
    3. α particle
    4. Neutron
    5. γ ray

46. When chromium metal is used to form K₂Cr₂O₇, the oxidation state of chromium changes from

    1. 0 to 4
    2. 3 to 6
    3. 2 to 6
    4. 0 to 6
    5. 2 to 4

47. Electron density studies have revealed that X and Y have an equal number of electrons. Which of the following could X and Y be?

    1. Ca⁺ and K
    2. H⁺ and He
    3. Cl and F
    4. O^– and S⁺
    5. None of the above

48. All halogens have similar reactivity because

    1. they have the same number of protons
    2. they have the same number of electrons
    3. they have similar outer shell electron configurations
    4. they have valence electrons with the same quantum numbers
    5. they have the same number of neutrons

49. K⁺ and Cl^– have the same

    1. atomic weight
    2. electronic configuration
    3. ionization potential
    4. number of protons and neutrons
    5. atomic radius

50. Which of the following has the largest ionic radius?

    1. Na⁺
    2. K⁺
    3. Mg⁺⁺
    4. Al³⁺
    5. Cl^–

51. When the following reaction is balanced, what is the net ionic charge on the right side of the equation?

    …H⁺ + …MnO₄^– + …Fe²⁺ → …Mn²⁺ + …Fe³⁺ + …H₂O
    1. +5
    2. +7
    3. +10
    4. +17
    5. The net ionic charge on either side must be zero.

52. Questions 52–54 refer to the following equation.

53. …Ag(NH₃)₂⁺ → …Ag⁺ + …NH₃

54. 52. What is the sum of the coefficients once the equation is balanced?

        1. 1
        2. 2
        3. 3
        4. 4
        5. 5

    53. How many moles of Ag(NH₃)²+ are required to produce 11 moles of ammonia?

        1. 1
        2. 2
        3. 5.5
        4. 11
        5. 22

    54. What is the percent composition by weight of Ag in Ag(NH₃)⁺₂?

        1. 4.2
        2. 19.7
        3. 76.1
        4. 80.3
        5. 95.8

55. If 88 g of C₃H₈ and 160 g of O₂ are allowed to react maximally to form CO₂ and H₂O, how many grams of CO₂ will be formed?

    1. 33
    2. 66
    3. 132
    4. 264
    5. None of the above

56. A 200 mL flask contains oxygen at 200 mm Hg, and a 300 mL flask contains neon at 100 mm Hg. The two flasks are connected so that each gas fills their combined volumes. Assuming no change in temperature, what is the partial pressure of neon in the final mixture?

    1. 60 mm Hg
    2. 80 mm Hg
    3. 100 mm Hg
    4. 150 mm Hg
    5. 200 mm Hg

57. What is the value of Z in the beta decay reaction below?

    
    1. 25
    2. 26
    3. 27
    4. 28
    5. 29

58. Questions 58–60 refer to the following experimental setups.

59. 

60. 58. Which of the following experimental setups will complete the circuit?

        1. II only
        2. III only
        3. III and IV
        4. II, III, and IV
        5. I, II, III, and IV

    59. What types of bonds are found in solid NaCl?

        1. Van der Waals
        2. Ionic
        3. Covalent
        4. Hydrogen
        5. Hydrophobic

    60. What would happen if the electrodes were put in a saturated solution of glucose dissolved in water?

        1. Light bulb would glow.
        2. Light bulb would remain dark.
        3. Apparatus would combust.
        4. Glucose molecules would dissociate.
        5. None of the above

61. Which of the following statements correctly characterizes a galvanic cell?

    1. Oxidation occurs at the anode, which is negative.
    2. Oxidation occurs at the anode, which is positive.
    3. Reduction occurs at the cathode, which is positive.
    1. II only
    2. III only
    3. I and III
    4. I, II, and III
    5. II and III

62. The heat of combustion of gaseous ammonia is 81 kcal/mole. How much heat is evolved in the reaction of 34 grams of ammonia with excess oxygen?

    1. 40.5 kcal
    2. 60.3 kcal
    3. 75.8 kcal
    4. 81 kcal
    5. 162 kcal

63. When there are two electrons in the 3s sublevel,

    1. they must occupy different orbitals
    2. the Heisenberg uncertainty principle predicts that they must periodically jump to the 3p sublevel
    3. the oxidation state of the atom must be 2+
    4. they must have opposite spins
    5. they are oppositely charged

64. In order to make a buffer solution, a weak monoprotic acid could be added to

    1. another acid
    2. another base
    3. its conjugate base
    4. its conjugate acid
    5. a strong base

65. How will the equilibrium of the following reaction be affected if the temperature is increased?

    N₂ (g) + 3H₂ (g) ⇌ 2NH₃ (g) ∆H = –30 kJ/mol
    1. It will be shifted to the right.
    2. It will be shifted to the left.
    3. It will be unaffected.
    4. The effect on the equilibrium cannot be determined without more information.
    5. None of the above

66. For a sample of an ideal gas of fixed weight and at a fixed temperature,

    1. the volume varies directly with the pressure exerted on it
    2. the volume varies inversely with the pressure exerted on it
    3. the pressure varies directly with the density of the gas
    1. I only
    2. II only
    3. III only
    4. I and II
    5. II and III

67. When 200 g of MgCl₂ is added to 1 kg of water, what is the molality of the solution?

    1. 200/(24 + 71)
    2. 0.200/(24 + 71)
    3. 0.2(24 + 71)
    4. 0.200
    5. 200

68. How many moles of water are formed by a mixture of 100 grams of H₂ and 100 grams of O₂? (Assume the reaction goes to completion.)

    1. 100/32 + 100/64
    2. 100 + 2(100/32)
    3. 2(100/32)
    4. 100(100/32)
    5. 200(100/32)

69. An oxide of arsenic contains 65.2% arsenic by weight. What is its simplest formula?

    1. AsO
    2. As2O₃
    3. AsO₂
    4. As₂O₅
    5. As₂O

If you finish before time is called, you may check your work on this section only. Do not turn to any other section in the test.