Chem 106 Final Exam Study Questions – Spring 2018

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Chem 106 Final Exam Study Questions – Spring 2018 Note: The answer key on the back contains chapter/sections from the text.

1. A student finds that the weight of an empty beaker is 12.024 g. She places a solid in the beaker to give a combined mass of 12.108 g. To how many significant figures is the mass of the solid known? A) 1 B) 2 C) 3 D) 4 E) 5

2. 8.2 milliseconds is equal to how many seconds? A) 8.2 103 s B) 8.2 102 s C) 8.2 10–3 s D) 8.2 10–2 s E) 0.82 s

3. Convert 348.4 m to decimeters. A) 3.484 104 dm B) 34.84 dm C) 3.484 dm D) 3.484 103 dm E) none of these

4. 473.1oF is equivalent to A) 280.6oC B) 245.1oC C) 794.0oC D) 909.2oC E) 262.8oC

5. Convert: –49.5oF = _______________ oC. A) –45.3oC B) –9.7oC C) –146.7oC D) –31.5oC E) 223.5oC

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6. The number 2.00152 rounded to four significant figures is

A) 2.002 B) 2.001 C) 2.000 D) 2.152 E) none of these

7. A graduated cylinder contains 20.0 mL of water. An irregularly shaped object is placed in the cylinder, and the water level rises to the 31.2-mL mark. If the object has a mass of 91.1 g, what is its density? A) 8.13 g/mL B) 0.123 g/mL C) 2.92 g/mL D) 4.56 g/mL E) none of these

8. An object is 148.7 inches in height. Express this height in centimeters. A) 58.54 cm B) 0.01708 cm C) 12.39 cm D) 377.7 cm E) 151.2 cm

9. Which of the following is a chemical change? A) Water condenses on a mirror. B) A damp towel dries. C) Peanuts are crushed. D) A “tin” can rusts. E) At least two of the above (a-d) exhibit a chemical change.

10. Which of the following is an element? A) air B) water C) salt D) helium E) sugar

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11. Which is an example of a homogeneous mixture? A) vodka B) oily water C) soil (dust) D) sodium chloride E) aluminum

12. How many protons, electrons, and neutrons does the isotope 57 326 Fe have?

13. Which pair have approximately the same mass? A) a hydrogen, 1

1H , and a deuterium, 21H , atom

B) a neutron and an electron C) a proton and a neutron D) an electron and a proton

14. Which of the following elements is an alkaline earth metal? A) Ca B) Cu C) Fe D) Na E) Sc

15. The total number of atoms indicated by the formula Ca3(PO4)3 is A) 6 B) 10 C) 16 D) 18 E) 7

16. Which of the following is an incorrect name for an acid? A) hydrocarbonate acid B) hydrocyanic acid C) acetic acid D) phosphoric acid E) sulfurous acid

17. Give the formula for calcium hydrogen carbonate.

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18. Write the correct formula for dinitrogen pentoxide.

19. The name for NO3- is ______________.

20. Give the formula for carbon monoxide.

21. Give the formula for hypochlorous acid.

22. The name for (NH4)2SO4 is ______________.

23. The name for 2 3 2C H O is _______________.

24. Give the formula for hydrosulfuric acid.

25. The name for (NH4)2CO3 is _______________.

26. Give the formula for chromium(III) iodide.

27. The name for Zn(OH)2 is ______________.

28. The binary compound PCl3 is called A) phosphorus chloride B) triphosphorus chloride C) monophosphorus trichloride D) phosphorus trichloride E) none of these

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29. When the following equation is balanced using the smallest possible integers, what is the number in front of the substance in bold type? P4O10 + H2O H3PO4 A) 10 B) 6 C) 4 D) 2 E) 1

30. Balance the equation H2O2(l) H2O(l) + O2(g)

Use the following to answer questions 31-34: Use the following choices to classify each reaction given below (more than one choice may apply). a. oxidation-reduction b. acid-base c. precipitation

31. HNO3(aq) + NaOH(aq) H2O(l) + NaNO3(aq)

32. HC2H3O2(aq) + CsOH(aq) H2O(l) + CsC2H3O2(aq)

33. 2HCl(aq) + Pb(OH)2(aq) PbCl2(s) + 2H2O(l)

34. Zr(s) + O2(g) ZrO2(s) Use the following to answer questions 35-37: Use the following choices to classify each reaction given below (more than one choice may apply). a. oxidation-reduction b. combustion c. synthesis d. decomposition

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35. 2Na(s) + H2(g) 2NaH(s)

36. C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g)

37. 2GaN(s) 2Ga(s) + N2(g)

38. True or false? The formula of a compound that expresses the smallest whole-number ratio of the atoms present is called the empirical formula. A) True B) False

39. Which represents the greatest number of atoms? A) 50.0 g Al B) 50.0 g Cu C) 50.0 g Zn D) 50.0 g Fe E) all the same

40. How many atoms of calcium are present in 52.5 g of calcium? A) 2.18 10–24 B) 3.16 1025 C) 6.02 1023 D) 7.89 1023 E) none of these

41. Calculate the mass of 15.6 moles of He. A) 19.6 B) 3.90 C) 9.39 1024 D) 62.4 E) 6.02 1023

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42. How many moles of Ca atoms are in 631.1 g Ca? A) 6.351 10–2 mol B) 15.75 mol C) 2.529 104 mol D) 31.49 mol E) 631.1 mol

43. Calculate the mass of 2.90 1019 molecules of HCl. A) 1.76 10–3 g B) 4.82 10–5 g C) 1.75 1043 g D) 1.32 10–6 g E) none of these

44. A 27.6-mol sample of Co represents how many atoms? A) 4.58 10–23 atoms B) 2.18 1022 atoms C) 9.79 1026 atoms D) 6.72 102 atoms E) 1.66 1025 atoms

45. 11.2 g of Cu represents how many moles? A) 7.12 102 mol B) 5.67 mol C) 0.176 mol D) 6.74 1024 mol E) none of these

46. What is the molar mass of nitroglycerin, C3H5(NO3)3? A) 179 g/mol B) 227 g/mol C) 199 g/mol D) 185 g/mol E) none of these

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47. Calculate the mass of 0.569 mol of H2SO4. A) 1.72 102 g B) 55.8 g C) 5.80 10–3 g D) 3.43 1023 g E) 3.36 1025 g

48. A gaseous compound containing carbon and hydrogen was analyzed and found to consist of 83.65% carbon by mass. What is the empirical formula of the compound? A) CH2 B) CH3 C) C7H16 D) CH E) C3H7

49. Determine the percentage composition (by mass) of oxygen in H2SO4. A) 16.31 % B) 31.95 % C) 65.25 % D) 23.48 % E) 78.25 %

50. Choose the pair of compounds with the same empirical formula. A) C2H2 and C6H6 B) NaHCO3 and Na2CO3 C) K2CrO4 and K2Cr2O7 D) H2O and H2O2

51. A compound contains 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen (by mass). Calculate the empirical formula. A) CH2O B) C2H2O C) CH4O D) C3H6O3 E) C2HO2

52. A compound has 40.68% carbon, 5.12% hydrogen, and 54.20% oxygen (by mass). Calculate its empirical formula.

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53. A 5.4-mol sample of KClO3 was decomposed according to the equation 3 22KClO (s) 2KCl(s) 3O (g)

How many moles of O2 are formed assuming 100% yield? A) 3.6 mol B) 4.5 mol C) 5.4 mol D) 2.7 mol E) 8.1 mol

54. Refer to the following unbalanced equation: C6H14 + O2 CO2 + H2O 1. Balance the equation 2. What mass of carbon dioxide (CO2) can be produced from 24.9 g of C6H14 and excess oxygen? A) 6.58 103 g B) 12.7 g C) 76.3 g D) 38.2 g E) 0.289 g

55. Consider the reaction 2 2 34Fe(s) 3O (g) 2Fe O (s)

If 11.8 g of iron(III) oxide (rust) is produced from a certain amount of iron, how many grams of oxygen are needed for this reaction? A) 7.09 g B) 3.55 g C) 1.58 g D) 2.36 g E) none of these

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56. For the reaction 3 2 2 2CaCO (s) 2HCl(aq) CaCl (aq) CO (g) H O(l)

how many grams of CaCl2 can be obtained if 42.8 g HCl is allowed to react with excess CaCO3? A) 130 g CaCl2 B) 261 g CaCl2 C) 0.587 g CaCl2 D) 65.1 g CaCl2 E) none of these

57. True or false? The frequency of the wave is the distance between two consecutive wave peaks. A) True B) False

58. The form of EMR that has less energy per photon than infrared rays but more energy per photon than radio waves is A) microwaves B) untraviolet C) gamma rays D) X rays E) none of these

59. Which color of visible light has the least amount of energy per photon? A) violet B) blue C) green D) yellow E) red

60. As the principal energy level increases in an atom's orbitals, the average distance of an electron energy level from the nucleus ______________. A) increases B) decreases C) stays the same D) varies E) none of these

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61. A given set of f orbitals consists of ______________ orbital(s). A) 1 B) 3 C) 5 D) 7 E) 9

62. The maximum electron capacity of an f sublevel is A) 18 B) 14 C) 10 D) 6 E) 2

63. The maximum number of electrons allowed in the p sublevel of the third principal level is A) 1 B) 2 C) 3 D) 6 E) 8

64. The maximum number of electrons allowed in the fourth energy level is A) 2 B) 4 C) 8 D) 18 E) 32

65. The number of d orbitals in the second principal energy level is A) 2 B) 6 C) 10 D) 14 E) none of these

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66. The number of unpaired electrons in a nitrogen atom is A) 1 B) 2 C) 3 D) 4 E) 5

67. Choose the correct ground-state electron configuration for oxygen. A) [He]2s22p4 B) [He]2s22p6 C) [Ne]2s22p4 D) [He]3s23p4 E) [He]1s22p6

68. The alkaline earth metals have how many valence electrons? A) 8 B) 7 C) 3 D) 2 E) 1

69. When moving down a group (family) in the periodic table, the number of valence electrons A) remains constant B) increases by 2 then 8 then 18 then 32 C) doubles with each move D) decreases regularly E) changes in an unpredictable manner

70. Which of the following atoms has the electron configuration 1s22s22p63s23p64s23d1? A) Sc B) Ca C) Sr D) Ar E) none of these

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71. Which electron configuration indicates a transition element? A) 1s22s22p63s13p6 B) 1s22s22p63s23p64s23d3 C) 1s22s22p5 D) 1s22s22p63s23p64s23d104p2 E) none of these

72. How many of the following electron configurations for the species in their ground state are correct? I. Ca: 1s22s22p63s23p64s2

II. Mg: 1s22s22p63s1

III. V: [Ar] 3s23d3

IV. As: [Ar] 4s23d104p3

V. P: 1s22s22p63p5 A) 1 B) 2 C) 3 D) 4 E) 5

73. Write the electron configuration for Cd.

74. Write the electron configuration for Cl.

75. Which of the following is ranked in order of largest to smallest atomic radius? A) Rb > Mn > S > Ge > F B) F > S > Ge > Mn > Rb C) Mn > Rb > F > S > Ge D) Rb > Ge > Mn > F > S E) Rb > Mn > Ge > S > F

76. Which of the following atoms has the highest ionization energy? A) Al B) Si C) P D) As E) Sb

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77. True or false: Covalent bonding occurs when electrons are shared by nuclei. A) True B) False

78. True or false? The greater the difference in electronegativity between two bonded atoms, the more polar the bond. A) True B) False

79. True or false? N2 is an example of a covalent bond. A) True B) False

80. The most electronegative element of those listed is A) Rb B) Cs C) Fr D) K E) Li

81. Which of the following has nonpolar bonds? A) H2S B) HCl C) Br2 D) OF2 E) All are nonpolar.

82. The number of polar covalent bonds in NH3 is A) 1 B) 2 C) 3 D) 4 E) none of these

83. Draw the Lewis electron structure for the sulfur atom.

84. Draw the Lewis structure for SiH4.

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85. Draw the Lewis structure for CO.

86. Which of the following has a double bond? A) H2O B) NH3 C) O2 D) CO E) H2S

Use the following to answer questions 87-88: Use the following choices to describe the molecular structure of each of the following molecules or ions. a. linear b. trigonal planar c. tetrahedral d. trigonal pyramid e. Bent (V-shaped)

87. CH4

88. PF3

89. Convert 2.8 102 atm to torr. A) 2.8 104 torr B) 0.37 torr C) 19 torr D) 4.1 104 torr E) 2.1 105 torr

90. The air in the inner tube of the tire of a racing bike has a pressure of 112.2 psi. Convert this pressure to atm. A) 0.1476 atm B) 7.633 atm C) 1649 atm D) 1.108 atm E) 112.2 atm

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91. Consider a gas at 1.00 atm in a 5.00-L container at 20.oC. What pressure does the gas exert when transferred to a volume of 3.01 L at 43oC? A) 3.57 atm B) 1.54 atm C) 0.371 atm D) 1.79 atm E) 0.649 atm

92. When analyzing ideal gases, the temperature must be measured on the Kelvin scale A) because otherwise you could calculate a negative volume. B) so that you are using an absolute scale. C) to directly measure the average kinetic energy of the gas particles. D) Both a and b are correct. E) a, b, and c are correct.

93. Determine the pressure exerted by 4.93 mol of gas in a 2.92-L container at 32oC. A) 4.43 atm B) 123.4 atm C) 42.3 atm D) 12.95 atm E) 37.8 atm

94. A sample of an ideal gas containing 0.707 mol is collected at 742 torr pressure and 31oC. Calculate the volume. A) 18.1 L B) 1.84 L C) 2.38 10–2 L D) 2.42 10–3 L E) none of these

95. 3.82 mol of CO2 at STP will occupy A) 93.4 L B) 85.6 L C) 1.13 10–1 L D) 42.8 L E) 85.6 g

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96. A vessel with an internal volume of 16.4 L contains 2.80 g of nitrogen gas, 0.403 g of hydrogen gas, and 79.9 g of argon gas. At 25oC, what is the pressure (in atm) inside the vessel? A) 124 atm B) 0.29 atm C) 922 atm D) 3.43 atm E) 0.292 atm

97. What volume of HCl(g) measured at STP can be produced from 3.97 g of H2 and excess Cl2 according to the following equation? 2 2H (g) Cl (g) 2HCl(g)

A) 44.1 L B) 88.2 L C) 178 L D) 359 L E) 133.4 L

98. At 1 atm of pressure and a temperature of 0°C, which phase(s) of H2O can exist? A) ice and water B) ice and water vapor C) water only D) water vapor only E) ice only

99. The normal boiling point of water is A) 0°F B) 32°F C) 273 K D) 373 K E) none of these

100. True or false? The bonding forces that hold the atoms of a molecule together are called intermolecular forces, whereas the forces that occur among molecules that cause them to aggregate to form a solid or a liquid are called intramolecular forces. A) True B) False

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101. The bonds between hydrogen and oxygen in a water molecule can be characterized as ______________. A) hydrogen bonds B) London forces C) intermolecular forces D) intramolecular forces E) dispersion forces

102. When a water molecule forms a hydrogen bond with another water molecule, which atoms are involved in the interaction? A) a hydrogen from one molecule and a hydrogen from the other molecule B) a hydrogen from one molecule and an oxygen from the other molecule C) an oxygen from one molecule and an oxygen from the other molecule D) two hydrogens from one molecule and one oxygen from the other molecule E) two hydrogens from one molecule and one hydrogen from the other molecule

103. Which of the following should have the lowest boiling point? A) CH4 B) C2H6 C) C3H8 D) C4H10 E) C5H12

Use the following to answer questions 104-108: Identify the major attractive force between particles in each of the compounds in questions as one the following: dipole-dipole hydrogen bonding London dispersion ionic

104. O2

105. Octane (C8H18)

106. CH3CH2OH

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107. Sodium chloride

108. CO

109. Consider the following compounds: CO NH3 CO2 CH4 H2

How many of the compounds above exhibit London dispersion forces? A) 1 B) 2 C) 3 D) 4 E) 5

110. Rank the following compounds from lowest to highest boiling point. CH3OH CH4 H2O C2H6 A) H2O < CH3OH < C2H6 < CH4 B) C2H6 < CH4 < CH3OH < H2O C) CH4 < C2H6 < CH3OH < H2O D) CH4 < C2H6 < H2O < CH3OH E) CH4 < CH3OH < C2H6 < H2O

111. In an atomic solid, what are the individual components? A) bonds B) atoms C) molecules D) ions E) electrons

Use the following to answer questions 112-117: Identify the type of solid as one of the following: Ionic Solid Molecular Solid Atomic Solid

112. dry ice (solid carbon dioxide)

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113. xenon

114. copper metal

115. SF6

116. sulfur (S8)

117. NaHCO3

118. The total mass of a solution is 167.7 g. The solvent mass is 125.2 g. What is the mass percent of the solute? A) 74.7 % B) 36.3 % C) 25.3 % D) 17.0 % E) Not enough information is given.

119. A 106.4-g sample of nitric acid solution that is 70.0% HNO3 (by mass) contains A) 74.5 mol HNO3 B) 1.18 mol HNO3 C) 1.69 mol HNO3 D) 4.69 103 mol HNO3 E) none of these

120. Determine the concentration of a solution made by dissolving 32.6 g of sodium chloride in 750.0 mL of solution. A) 0.418 M B) 43.5 M C) 0.558 M D) 0.744 M E) 24.5 M

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121. A solution is prepared by dissolving 5.88 g of Na2SO4 in enough water to make 225 mL of solution. Calculate the solution molarity. A) 0.0414 M B) 1.32 M C) 0.528 M D) 0.184 M E) 0.409 M

122. How many grams of CaCl2 (molar mass = 111.0 g/mol) are needed to prepare 3.02 L of 0.500 M CaCl2 solution? A) 180 g B) 168 g C) 193 g D) 291 g E) 218 g

123. What mass of calcium chloride, CaCl2, is needed to prepare 2.760 L of a 1.56 M solution? A) 4.31 g B) 325 g C) 196 g D) 62.7 g E) 478 g

124. A 89.15-g sample of NaCl is dissolved in 250.0 mL of solution. Calculate the molarity of this solution. A) 1.525 M B) 356.6 M C) 6.102 M D) 7.335 M E) none of these

125. What volume of 12.0 M nitric acid is required to prepare 3.19 L of 0.100 M nitric acid? A) 0.376 L B) 37.6 L C) 0.319 L D) 0.0266 L E) 3.76 L

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126. A 58.50-g sample of Ba(OH)2 is dissolved in enough water to make 1.20 L of solution. How many milliliters of this solution must be diluted with water in order to make 1.00 L of 0.100 M Ba(OH)2? (Ignore significant figures for this problem.) A) 351 mL B) 70.2 mL C) 341 mL D) 285 mL E) 585 mL

127. Assume that vinegar is a 0.852 M solution of acetic acid (HC2H3O2) in water. What volume of 0.2136 M NaOH would be needed to completely neutralize 5.18 mL of vinegar? A) 4.41 mL B) 1.11 mL C) 1.30 mL D) 20.7 mL E) 4.00 mL

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Answer Key

1. B Chapter: Ch 2.5

2. C Chapter: Ch 2.2,2.6

3. D Chapter: Ch 2.2,2.6


5. A Chapter: Ch 2.7



8. D Chapter: Ch 2.6




12. The isotope contains 26 protons, 23 electrons, and 31 neutrons. Chapter: Ch 4.7,4.10

13. C Chapter: Ch 4.6




17. Ca(HCO3)2 Chapter: Ch 5.5

18. N2O5 Chapter: Ch 5.3

19. nitrate ion Chapter: Ch 5.5

20. CO Chapter: Ch 5.3

21. HClO Chapter: Ch 5.6

22. ammonium sulfate

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Chapter: Ch 5.5 23. acetate ion

Chapter: Ch 5.5 24. H2S

Chapter: Ch 5.6 25. ammonium carbonate

Chapter: Ch 5.5 26. CrI3

Chapter: Ch 5.2 27. zinc hydroxide

Chapter: Ch 5.5 28. D

Chapter: Ch 5.3 29. E

Chapter: Ch 6.3 30. 2H2O2(l) 2H2O(l) + O2(g)

Chapter: Ch 6.3 31. b

Chapter: Ch 7.6 32. b

Chapter: Ch 7.6 33. b; c

Chapter: Ch 7.6 34. a

Chapter: Ch 7.7 35. a; c

Chapter: Ch 7.7 36. a; b

Chapter: Ch 7.7 37. a; d

Chapter: Ch 7.7 38. A




Chapter: Ch 8.3 42. B



Chapter: Ch 8.3

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48. E Chapter: Ch 8.8




52. C2H3O2 Chapter: Ch 8.8


54. C Chapter: Ch 6.3,9.3













67. A

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Chapter: Ch 11.10 73. [Kr] 5s24d10

Chapter: Ch 11.10 74. [Ne] 3s23p5


Chapter: Ch 11.11 76. C







Chapter: Ch 12.3 83.

Chapter: Ch 12.6

84.

Chapter: Ch 12.6

85.

Chapter: Ch 12.7

S . .

.

. .

.

H — Si — H

H

—

—

H

C O. .

.

. — — —

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87. c Chapter: Ch 12.9

88. d Chapter: Ch 12.9
















104. London dispersion Chapter: Ch 14.3

105. London dispersion Chapter: Ch 14.3

106. hydrogen bonding Chapter: 14.3

107. ionic Chapter: Ch 12.5

108. dipole-dipole

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Chapter: Ch 14.5,14.6 112. molecular solid

Chapter: Ch 14.5,14.6 113. atomic solid

Chapter: Ch 14.5,14.6 114. atomic solid



Chapter: Ch 14.5,14.6 117. ionic solid

Chapter: Ch 14.5,14.6 118. C










Chapter: Ch 15.7

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Chem 106 Final Exam Study Questions – Spring 2018