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

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1 Chem 106 Final Exam Study Questions Summer 2017 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 g. She places a solid in the beaker to give a combined mass of g. To how many significant figures is the mass of the solid known? A) 1 B) 2 C) 3 D) 4 E) milliseconds are equal to how many seconds? A) s B) s C) s D) s E) 0.15 s 3. Convert m to decimeters. A) dm B) dm C) dm D) dm o F is equivalent to A) o C B) o C C) o C D) o C E) o C 5. Convert: 43.5 o F = o C. A) 41.9 o C B) 6.4 o C C) o C D) 20.7 o C E) o C Page 1

2 6. The number rounded to four significant figures is A) B) C) D) 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 60.0 g, what is its density? A) 5.36 g/ml B) g/ml C) 1.92 g/ml D) 3.00 g/ml 8. An object is inches in height. Express this height in centimeters. A) cm B) cm C) cm D) cm E) 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 Page 2

3 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 Fe have? 13. Which pair have approximately the same mass? A) a hydrogen, 1 H, and a deuterium, 2 H, atom 1 1 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) 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. Page 3

4 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 CHO 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 Page 4

5 29. When the following equation is balanced using the smallest possible integers, what is the number in front of the substance in bold type? P 4 O 10 + H2O H3PO4 A) 10 B) 6 C) 4 D) 2 E) 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 Page 5

6 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 53.0 g of calcium? A) B) C) D) Calculate the mass of 22.4 moles of He. A) 26.4 B) 5.60 C) D) 89.7 E) Page 6

7 42. How many moles of Ca atoms are in g Ca? A) mol B) mol C) mol D) mol E) mol 43. Calculate the mass of molecules of HCl. A) g B) g C) g D) g 44. A 42.2-mol sample of Co represents how many atoms? A) atoms B) atoms C) atoms D) atoms E) atoms g of Cu represents how many moles? A) mol B) 4.64 mol C) mol D) mol 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 Page 7

8 47. Calculate the mass of mol of H2SO4. A) g B) 20.8 g C) g D) g E) 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) % B) % C) % D) % E) % 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. Page 8

9 53. A 5.8-mol sample of KClO3 was decomposed according to the equation 2KClO 3(s) 2KCl(s) 3O 2(g) How many moles of O2 are formed assuming 100% yield? A) 3.9 mol B) 4.8 mol C) 5.8 mol D) 2.9 mol E) 8.7 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 18.3 g of C6H14 and excess oxygen? A) g B) 9.3 g C) 56.1 g D) 28.0 g E) g 55. Consider the reaction 2Fe(s) 3O 2(g) Fe2O 3(s) If 11.4 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) 3.43 g B) 6.85 g C) 1.52 g D) 2.28 g Page 9

10 56. For the reaction CaCO 3(s) 2HCl(aq) CaCl 2(aq) CO 2(g) H2O(l) how many grams of CaCl2 can be obtained if 31.2 g HCl is allowed to react with excess CaCO3? A) 95 g CaCl2 B) 190 g CaCl2 C) g CaCl2 D) 47.5 g CaCl2 57. For the reaction 2S(s) 3O 2(g) 2SO 3(g) if 6.28 g of S is reacted with 10.0 g of O2, how many grams of SO3 will be produced? A) 31.4 g B) 7.84 g C) 16.7 g D) 15.7 g 58. Sodium and water react according to the equation 2Na(s) 2H2O(l) 2NaOH(aq) H 2(g) What number of moles of H2 will be produced when 4.0 mol Na is added to 3.5 mol H2O? A) 1.8 mol B) 7.0 mol C) 2.0 mol D) 3.5 mol E) 8.0 mol Page 10

11 59. Consider the following unbalanced equation: C2H5OH(g) + O2(g) CO2(g) + H2O(l) If 1.92 g of ethanol reacts with 13.8 g of oxygen, how many moles of water are produced? A) mol B) mol C) mol D) mol E) mol 60. True or false? The frequency of the wave is the distance between two consecutive wave peaks. A) True B) False 61. 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 62. Which color of visible light has the least amount of energy per photon? A) violet B) blue C) green D) yellow E) red 63. 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 Page 11

12 64. A given set of f orbitals consists of orbital(s). A) 1 B) 3 C) 5 D) 7 E) The maximum electron capacity of an f sublevel is A) 18 B) 14 C) 10 D) 6 E) 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) The maximum number of electrons allowed in the fourth energy level is A) 2 B) 4 C) 8 D) 18 E) The number of d orbitals in the second principal energy level is A) 2 B) 6 C) 10 D) 14 Page 12

13 69. The number of unpaired electrons in a nitrogen atom is A) 1 B) 2 C) 3 D) 4 E) Choose the correct ground-state electron configuration for oxygen. A) [He]2s 2 2p 4 B) [He]2s 2 2p 6 C) [Ne]2s 2 2p 4 D) [He]3s 2 3p 4 E) [He]1s 2 2p The alkaline earth metals have how many valence electrons? A) 8 B) 7 C) 3 D) 2 E) 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 73. Which of the following atoms has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1? A) Sc B) Ca C) Sr D) Ar Page 13

14 74. Which electron configuration indicates a transition element? A) 1s 2 2s 2 2p 6 3s 1 3p 6 B) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 C) 1s 2 2s 2 2p 5 D) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p How many of the following electron configurations for the species in their ground state are correct? I. Ca: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 II. Mg: 1s 2 2s 2 2p 6 3s 1 III. V: [Ar] 3s 2 3d 3 IV. As: [Ar] 4s 2 3d 10 4p 3 V. P: 1s 2 2s 2 2p 6 3p 5 A) 1 B) 2 C) 3 D) 4 E) Write the electron configuration for Cd. 77. Write the electron configuration for Cl. 78. 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 79. Which of the following atoms has the highest ionization energy? A) Al B) Si C) P D) As E) Sb Page 14

15 80. True or false: Covalent bonding occurs when electrons are shared by nuclei. A) True B) False 81. True or false? The greater the difference in electronegativity between two bonded atoms, the more polar the bond. A) True B) False 82. True or false? N2 is an example of a covalent bond. A) True B) False 83. The most electronegative element of those listed is A) Rb B) Cs C) Fr D) K E) Li 84. Which of the following has nonpolar bonds? A) H2S B) HCl C) Br2 D) OF2 E) All are nonpolar. 85. The number of polar covalent bonds in NH3 is A) 1 B) 2 C) 3 D) Draw the Lewis electron structure for the sulfur atom. 87. Draw the Lewis structure for SiH4. Page 15

16 88. Draw the Lewis structure for CO. 89. Which of the following has a double bond? A) H2O B) NH3 C) O2 D) CO E) H2S Use the following to answer questions 90-91: 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) 90. CH4 91. PF3 92. Convert atm to torr. A) torr B) 0.34 torr C) 18 torr D) torr E) torr 93. The air in the inner tube of the tire of a racing bike has a pressure of psi. Convert this pressure to atm. A) atm B) atm C) 1485 atm D) atm E) atm Page 16

17 94. Consider a gas at 1.00 atm in a 5.00-L container at 20. o C. What pressure does the gas exert when transferred to a volume of 2.41 L at 43 o C? A) 4.46 atm B) 1.92 atm C) atm D) 2.24 atm E) atm 95. 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. 96. Determine the pressure exerted by 4.50 mol of gas in a 2.92-L container at 32 o C. A) 4.05 atm B) atm C) 38.6 atm D) atm E) 34.5 atm 97. A sample of an ideal gas containing mol is collected at 742 torr pressure and 31 o C. Calculate the volume. A) 16.4 L B) 1.67 L C) L D) L mol of CO2 at STP will occupy A) 28.9 L B) 26.4 L C) L D) 13.2 L E) 26.4 g Page 17

18 99. A vessel with an internal volume of 16.8 L contains 2.80 g of nitrogen gas, g of hydrogen gas, and 79.9 g of argon gas. At 25 o C, what is the pressure (in atm) inside the vessel? A) 121 atm B) 0.28 atm C) 945 atm D) 3.35 atm E) atm 100. What volume of HCl(g) measured at STP can be produced from 2.69 g of H2 and excess Cl2 according to the following equation? H 2(g) Cl 2(g) 2HCl(g) A) 29.9 L B) 59.8 L C) 121 L D) 243 L E) 90.4 L 101. 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 102. The normal boiling point of water is A) 0 F B) 32 F C) 273 K D) 373 K 103. 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 Page 18

19 104. 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 105. 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 106. 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 : Identify the major attractive force in each of the following molecules O2 A) dipole-dipole B) London dispersion C) ionic D) hydrogen bonding 108. CO A) dipole-dipole B) London dispersion C) ionic D) hydrogen bonding Page 19

20 109. 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) 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. Name the type of crystalline solid formed by copper. A) molecular solid B) atomic solid C) ionic solid D) amorphous solid 112. The total mass of a solution is g. The solvent mass is g. What is the mass percent of the solute? A) 83.2 % B) 27.8 % C) 16.8 % D) 11.3 % E) Not enough information is given. Page 20

21 113. A g sample of nitric acid solution that is 70.0% HNO3 (by mass) contains A) 76.5 mol HNO3 B) 1.21 mol HNO3 C) 1.73 mol HNO3 D) mol HNO Determine the concentration of a solution made by dissolving 37.8 g of sodium chloride in ml of solution. A) M B) 50.4 M C) M D) M E) 28.4 M 115. A solution is prepared by dissolving 7.24 g of Na2SO4 in enough water to make 225 ml of solution. Calculate the solution molarity. A) M B) 1.63 M C) M D) M E) M 116. How many grams of CaCl2 (molar mass = g/mol) are needed to prepare 5.47 L of M CaCl2 solution? A) 316 g B) 304 g C) 329 g D) 427 g E) 354 g 117. What mass of calcium chloride, CaCl2, is needed to prepare L of a 1.56 M solution? A) 5.58 g B) 422 g C) 255 g D) 48.4 g E) 619 g Page 21

22 118. A g sample of NaCl is dissolved in ml of solution. Calculate the molarity of this solution. A) M B) M C) M D) M 119. What volume of 12.0 M nitric acid is required to prepare 3.05 L of M nitric acid? A) L B) 39.3 L C) L D) L E) 3.93 L 120. A 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 M Ba(OH)2? (Ignore significant figures for this problem.) A) 400 ml B) 61.7 ml C) 300 ml D) 250 ml E) 514 ml 121. Assume that vinegar is a M solution of acetic acid (HC2H3O2) in water. What volume of M NaOH would be needed to completely neutralize 5.90 ml of vinegar? A) 5.03 ml B) 1.26 ml C) 1.48 ml D) 23.5 ml E) 4.00 ml Page 22

23 Answer Key 1. B Chapter: Ch C Chapter: Ch 2.2, D Chapter: Ch 2.2, B Chapter: Ch A Chapter: Ch A Chapter: Ch A Chapter: Ch D Chapter: Ch D Chapter: Ch D Chapter: Ch A Chapter: Ch The isotope contains 26 protons 23 electrons 31 neutrons Chapter: Ch 4.7, C Chapter: Ch A Chapter: Ch D Chapter: Ch A Chapter: Ch Ca(HCO3)2 Chapter: Ch N2O5 Chapter: Ch nitrate ion Chapter: Ch CO Chapter: Ch HClO Chapter: Ch ammonium sulfate Chapter: Ch acetate ion Chapter: Ch H2S Chapter: Ch ammonium carbonate Chapter: Ch CrI3 Chapter: Ch zinc hydroxide Chapter: Ch D Chapter: Ch E Chapter: Ch H2O2(l) 2H2O(l) + O2(g) Chapter: Ch b Chapter: Ch b Chapter: Ch b; c Chapter: Ch a Chapter: Ch a; c Chapter: Ch a; b Chapter: Ch a; d Chapter: Ch A Chapter: Ch A Chapter: Ch D Chapter: Ch D Chapter: Ch B Chapter: Ch A Chapter: Ch 8.5 Page 23

24 44. E Chapter: Ch C Chapter: Ch B Chapter: Ch B Chapter: Ch E Chapter: Ch C Chapter: Ch A Chapter: Ch A Chapter: Ch C2H3O2 Chapter: Ch E Chapter: Ch C Chapter: Ch 6.3, B Chapter: Ch D Chapter: Ch D Chapter: Ch A Chapter: Ch B Chapter: Ch B Chapter: Ch A Chapter: Ch E Chapter: Ch A Chapter: Ch D Chapter: Ch B Chapter: Ch D Chapter: Ch E Chapter: Ch E Chapter: Ch C Chapter: Ch A Chapter: Ch D Chapter: Ch A Chapter: Ch A Chapter: Ch B Chapter: Ch B Chapter: Ch [Kr] 5s 2 4d 10 Chapter: Ch [Ne] 3s 2 3p 5 Chapter: Ch E Chapter: Ch C Chapter: Ch A Chapter: Ch A Chapter: Ch A Chapter: Ch E Chapter: Ch C Chapter: Ch C Chapter: Ch S.. Chapter: Ch 12.6 Page 2

25 H H Si H H Chapter: Ch C O. Chapter: Ch C Chapter: Ch c Chapter: Ch d Chapter: Ch E Chapter: Ch B Chapter: Ch D Chapter: Ch E Chapter: Ch C Chapter: Ch A Chapter: Ch B Chapter: Ch D Chapter: Ch B Chapter: Ch A Chapter: Ch D Chapter: Ch B Chapter: Ch D Chapter: Ch B Chapter: Ch A Chapter: Ch B Chapter: Ch A Chapter: Ch E Chapter: Ch C Chapter: Ch B Chapter: Ch 14.5, C Chapter: Ch B Chapter: Ch D Chapter: Ch D Chapter: Ch B Chapter: Ch E Chapter: Ch C Chapter: Ch D Chapter: Ch A Chapter: Ch D Chapter: Ch 15.7 Page 3

Note: The answer key on the back contains chapter/sections from the text.

Chem 106 Final Exam Study Questions Spring 2017 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