CHM 101 Sample Exam Answers (Last Page) Part 2

CHM 101 Sample Exam Questions 1. You are responsible for the information on this page. Please read it carefully. 2. Print your name and 10-digit affiliate identification number on the separate answer sheet. 3. Code your name and 10-digit affiliate identification number on the separate answer sheet. Use only a #2 pencil. 4. Do all calculations on the exam pages. Do not make any unnecessary marks on the answer sheet. 5. This exam consists of 25 multiple choice questions worth 4 points each and a periodic table. Make sure you have them all. 6. Choose the best answer to each of the questions and answer it on the computer-graded answer sheet. Read all responses before making a selection. 7. Bring your exam to class next period. Potentially useful information: 1 inch = 2.54 cm (exactly) 1 lb = 453.6 g 1 quart = 0.9464 L 4 quart = 1 gallon 9 º F = 32 + º C 5 5 º C = (º F - 32) 9 K = ºC + 273.15 Avogadro’s Number: 6.022 × 1023 E = hν (h = 6.626 × 10-34 J∙s) c = λν (c = 2.998 × 108 m/s) 109 nm = 1 m Activity Series Mg Al Zn Fe Ni Pb H2 Cu Hg Ag Au PV=nRT R=0.0821 (L·atm)/(mol·K) 760 mmHg = 760 torr = 1 atm P1 V1 P2 V2 = T1 T2 q = mC∆T (where q is heat, m is mass, C is specific heat, and ∆T is change in temperature) Chapter 1 1. A typical aspirin tablet may contain 325 mg of active ingredient. How many pounds of active ingredient does a typical aspirin tablet contain? A. 0.7164903 lb B. 1.39569 lb C. 7.16 × 10–4 lb D. 139 lb E. 716 lb 2. A lab instructor gave each of four groups of students a vial containing an unknown liquid. Each group ran four trials to determine the volume of the liquid. If the true volume is 156.2 mL which set of data is most precise yet not very accurate? A. B. C. D. E. Group I: 152.1 mL, 160.3 mL, 162.8 mL, 145.3 mL Group II: 168.7 mL, 168.6 mL, 168.7 mL, 168.8 mL Group III: 156.1 mL, 156.3 mL, 156.2 mL, 156.1 mL Group IV: 153.3 mL, 153.2 mL, 159.1 mL, 159.3 mL All of the above have the same precision. 3. Consider the following calculation: (8.32 × 10-5 ) × (0.45) (6.022 × 1023 ) How many significant figures should be reported in the answer? A. 1 B. 2 C. 3 D. 4 E. 23 4. Which of the following is considered an elemental substance? A. F2 B. BrCl3 C. Ba D. Al2O3 E. A and C are both elemental substances 5. What happens when water boils? A. B. C. D. E. A physical change occurs in which gaseous water molecules form. A physical change occurs in which hydrogen and oxygen molecules form. A chemical change occurs in which gaseous water molecules form. A chemical change occurs in which hydrogen and oxygen atoms form. Water boiling does not involve physical or chemical change. 6. Ethanol has a density of 0.789 g/mL. Which of the following liquids will float on ethanol? A. B. C. D. E. Mercury (density = 15.534 g/mL) Mineral oil (density = 0.86 g/mL) Water (density = 0.998 g/mL) Carbon tetrachloride (density = 1.589 g/mL) Ethanol will float on all of these. 7. Two students are working with two different bars of gold. One bar has a mass of 100 g and the other has a mass of 200 g. Which of the following quantities is the same for these two bars of gold? A. size B. mass C. density D. melting point E. Both C and D -2- 8. Which of these images is the best representation of a mixture of elements? A. B. C. D. E. 9. How many of the above boxes contain a diatomic element? A. 1 B. 2 C. 3 D. 4 E. 5 10. Which of the following quantities has the smallest mass? A. B. C. D. E. 4.2 × 10-3 kg 1.0 g 3.0 × 103 mg 600 mg 0.010 kg 11. A student working in a research lab finds a piece of metal with a volume of 125 cm3. If this metal has a mass of 337.5 g, which of the following is the metal most likely to be? A. B. C. D. E. aluminum (density = 2.70 g/cm3) sodium chloride (density = 2.16 g/cm3) lead (density = 11.3 g/cm3) gold (density = 19.3 g/cm3) The students discovered a new metal with a density of 0.370 g/cm3. Chapter 2 12. Which of these elements is described as a noble gas? A. barium B. fluorine C. sodium D. sulfur E. argon 13. Prior to Rutherford’s gold foil experiment, Thompson’s “plum pudding” model was consistent with experimental data of the time. Rutherford’s experiment showed that Thompson’s model was incorrect. Which of the following statements best describe Rutherford’s nuclear model of the atom? A. B. C. D. E. The heavy subatomic particles (protons and neutrons) reside in the nucleus. The light subatomic particles (protons and neutrons) reside in the nucleus. All subatomic particles have the same mass. Mass is spread uniformly throughout the atom. Both C and D. -3- 14. Dalton proposed his atomic theory in the early 1800s. Late in that century and into the beginning of the next, electrons and protons were discovered. Which of the following statements is not a postulate of Dalton’s atomic theory? A. B. C. D. E. All matter is made up of atoms. Atoms are composed of electrons, protons, and neutrons. Atoms of a particular element are all alike. Atoms of different elements are different from one another. Atoms can neither be created nor destroyed. 15. Which of the following isotopes has 45 neutrons? A. 80 35 Br B. 80 36 Kr C. 78 34 Se D. 34 17 Cl E. 103 45 Rh 16. Which of the following are the same for different isotopes of a particular element? A. B. C. D. E. neutrons protons protons and neutrons protons, neutrons, and electrons subatomic particles 17. Which combination of neutrons, protons, and electrons is correct for an ion with the symbol A. B. C. D. E. 63 29 Cu 2+ ? 29 neutrons, 34 protons, and 32 electrons 29 neutrons, 63 protons, and 65 electrons 29 neutrons, 34 protons, and 34 electrons 34 neutrons, 29 protons, and 31 electrons 34 neutrons, 29 protons, and 27 electrons Chapter 3 18. Which of the following combinations of elements will give an ionic compound? A. B. C. D. E. nitrogen and oxygen sulfur and oxygen calcium and oxygen phosphorus and chlorine xenon and chlorine 19. The image below represents a species having no charge and consisting of two nonmetals. How is this species best described? A. B. C. D. E. Ionic compound Homogeneous mixture Heterogeneous mixture Molecular compound Polyatomic ion -4- 20. The image below represents a compound containing sodium and chlorine. Which statement best describes this compound? A. B. C. D. E. The compound is molecular and consists of neutral sodium atoms and neutral chlorine atoms. The compound is ionic and consists of neutral sodium atoms and neutral chlorine atoms. The compound is ionic and consists of Na– ions and Cl+ ions. The compound is ionic and consists of Na+ ions and Cl– ions. This a homogeneous mixture containing Na atoms and Cl atoms. 21. What is the correct formula for the nitrate ion? D. NO2 A. N3– E. NO3 B. NO2– C. NO3– 22. Which of the following formulas with corresponding names is correct? A. B. C. D. E. Mn(NO3)2 Cr2(SO4)3 BaF PCl3 CO32– manganese(II) nitride chromium sulfate barium monofluoride phosphorous trichloride trioxocarbonate 23. What is the correct name for the compound with the formula Fe2(SO4)3? A. B. C. D. E. iron(III) sulfate diiron trisulfate iron sulfide iron sulfate iron(II) sulfate 24. What is the formula for a compound composed of sodium ions and phosphate ions? A. Na(PO4)3 B. Na3PO4 C. NaP D. NaP3 E. Na3P Chapter 4 25. What is the molar mass of chromium(III) nitrate, Cr(NO3)3? A. 82.00 g/mol B. 114.00 g/mol C. 210.00 g/mol D. 238.01 g/mol E. 6.022 × 1023 g/mol -5- 26. How many moles of calcium bromide (molar mass = 199.9 g/mol) are in 327 grams of this compound? A. 0.611 mol D. 6.54 × 104 mol B. 1.64 mol E. 9.85 × 1023 mol C. 2.73 mol 27. When 125 g of sodium carbonate, Na2CO3 (molar mass = 106.0 g/mol), is dissolved in water, how many sodium ions are present in the solution? A. 0.0128 Na+ ions B. 78.1 Na+ ions C. 1.30 × 10–23 Na+ ions D. 7.10 × 1023 Na+ ions E. 1.42 × 1024 Na+ ions 28. Ethylene glycol, used in automobile coolant, has the formula C2H6O2 (molar mass = 62.0 g/mol). What is the mass percent of oxygen in this compound? A. 9.68% B. 19.4% C. 25.8% D. 38.7% E. 51.6% 29. Which of the following is not an empirical formula? A. NO2 B. N2O4 C. N2O5 D. CO2 E. SO3 30. A compound containing carbon, hydrogen, and oxygen has the empirical formula CH2O. What is the molecular formula for this compound if it has a molar mass of 180.2 g/mol? A. CH2O B. C2H4O2 C. C3H6O3 D. C4H8O4 E. C6H12O6 Chapter 5 31. Magnesium reacts with hydrobromic acid to produce magnesium bromide and hydrogen gas. Which of the following equations best represents the reaction under normal conditions (room temperature, reactants and products in common form, atmospheric pressure, etc.)? A. B. C. D. E. Mg(l) + HBr(g) → MgBr(l) + H2(s) Mg(s) + 2HBr(aq) → MgBr2(aq) + H2(g) Mg(s) + HBr(l) → MgBr(aq) + H(g) Mg(s) + HBr(aq) → MgBr2(aq) + H2(g) Mg(s) + HBr(l) → MgBr(aq) + H2(g) 32. When the following equation is correctly balanced, what is the coefficient in front of water? ___C6H12(l) + ___O2(g) → ___CO2(g) + ___H2O(g) A. 6 B. 12 C. 16 D. 18 E. 34 -6- 33. Hydrogen gas reacts with nitrogen gas to form ammonia, NH3. Which of the following explanations best describes this reaction on a microscopic level? (Hint: Write a balanced equation for this reaction.) A. B. C. D. E. One hydrogen cation reacts with three nitrogen anions to form the ionic compound, ammonia. Three hydrogen cations reaction with one nitrogen anion to form the ionic compound, ammonia. Three hydrogen molecules react with one nitrogen molecule to form two ammonia molecules. Three hydrogen atoms react with one nitrogen atom to form one ammonia molecule. Three moles of hydrogen atoms react with 1 mole of nitrogen atoms to form one ammonia molecules. 34. Which of the following equations represents a single replacement reaction? A. B. C. D. E. 2H2O2(aq) → 2H2O(l) + O2(g) 2Ag(s) + Cu(NO3)2(aq) → 2AgNO3(aq) + Cu(s) CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) Pb(NO3)2(aq) + 2KCl(aq) → PbCl2(s) + 2KNO3(aq) 2Mg(s) + O2(l) → 2MgO(s) 35. Which of the following metals can be used to recover elemental silver from a silver nitrate solution in a single displacement reaction? A. copper B. lead C. gold D. zinc E. more than one of the above 36. Which of the following ionic compounds is most soluble in water? A. ZnS B. PbCl2 C. NaNO3 D. AgCl E. CaCO3 37. Which of the following equations best describes the expected reaction between potassium and bromine under normal conditions? A. B. C. D. E. K(s) + Br2(l) → KBr2(s) K2(s) + Br2(l) → K2Br2(s) 2K(s) + Br2(l) → 2KBr(s) P(g) + Br(g) → PBr(g) 2P(g) + 3Br2(g) → 2PBr3(s) 38. How is the reaction in the previous question best described? A. B. C. D. E. Synthesis (combination) Decomposition Single replacement Double replacement Combustion 39. When a solution of manganese(II) nitrate is added to a solution of sodium hydroxide, a solid forms as a result of a double displacement reaction. What is the net ionic equation for this reaction? A. B. C. D. E. Mn(NO3)2(aq) + 2NaOH(aq) → Mn(OH)2(s) + 2NaNO3(s) Mn(aq) + OH(aq) → MnOH(aq) Mn2+(aq) + 2OH–(aq) → Mn(OH)2(s) Na(aq) + NO3(aq) → NaNO3(s) Na+(aq) + NO3–(aq) → NaNO3(s) -7- 40. The spectator ion(s) for the reaction in the previous question is (are)? A. sodium and nitrate ions B. manganese and hydroxide ions C. manganese ions, only D. nitrate ions, only E. sodium ions, only 41. When solid sodium carbonate, Na2CO3, is put in hydrochloric acid, HCl, a gas forms. This reaction can be visualized as a double replacement reaction followed by a decomposition. What is the most likely identity of the gas that forms in the reaction? A. Na+ B. NaCl C. H2CO3 D. Cl– E. CO2 42. When sodium chlorate, NaClO3, is heated, sodium chloride and oxygen gas are produced. What type of reaction is this? A. B. C. D. E. decomposition single replacement double replacement combination (synthesis) precipitation 43. When sodium sulfate dissolves in water what are all the species present in the solution? A. B. C. D. E. Na atoms, S atoms, and O2 molecules Na atoms, S atoms, and O atoms Na atoms, S atoms, and H2O molecules Na+ ions, SO42– ions, and H2O molecules Na+ ions and SO42– ions, only Chapter 6 44. When 75.0 g of hot water was added to 150.0 g of water at 22.0 ºC, the final temperature was 35.0ºC. How much heat was gained by the cold water? (The specific heat of water is 4.184 J/g·ºC.) A. 54.4 J B. 4080 J C. 6900 J D. 8160 J E. 94,140 J 45. Suppose 6.0 mol of hydrogen is placed in a vessel containing 4.0 mol of nitrogen and allowed to react according to the equation: 3H2(g) + N2(g) → 2NH3(g) What will be present in the vessel after the reaction is complete? A. B. C. D. E. 4.0 mol NH3 only 4.0 mol NH3 and unreacted nitrogen 6.0 mol NH3 and unreacted hydrogen 8.0 mol NH3 and unreacted hydrogen 8.0 mol NH3 and unreacted nitrogen -8- 46. Hydrogen gas reacts with oxygen gas to produce water. The image below shows a mixture of molecules before the reaction occurs. What is present when the reaction is complete? (The larger molecules represent oxygen.) A. B. C. D. E. three water molecules and one hydrogen molecule three water molecules and seven hydrogen molecules six water molecules and one hydrogen molecule six water molecules and one oxygen molecule seven water molecules and one oxygen molecule 47. Magnesium metal reacts with oxygen gas to form magnesium oxide. How many moles of magnesium oxide are expected from the complete reaction of 0.0376 mol of O2? (Hint: Write a balanced equation for the reaction.) A. 0.0376 mol MgO B. 0.0188 mol MgO C. 0.0752 mol MgO D. 1.20 mol MgO E. 1.52 mol MgO 48. Sodium metal reacts with liquid bromine to form sodium bromide. How many grams of bromine are required to react with 75.0 g of sodium? 2Na(s) + Br2(l) → 2NaBr(s) A. 3.26 g Br2 B. 150 g Br2 C. 261 g Br2 D. 521 g Br2 E. 1040 g Br2 49. When 20.0 g of sodium is placed in 100.0 g of bromine a reaction occurs that is described in the previous question. How many moles of sodium react? A. 0.313 mol Na B. 0.626 mol Na C. 0.870 mol Na D. 1.25 mol Na E. 1.73 mol Na 50. Sodium metal reacts with liquid bromine to form sodium bromide. How many grams of bromine are required to react with 75.0 g of sodium? 2Na(s) + Br2(l) → 2NaBr(s) A. 3.26 g Br2 B. 150 g Br2 C. 261 g Br2 D. 521 g Br2 E. 1040 g Br2 51. How many joules of heat energy are required to heat 375 g of iron from 25.0ºC to 200.0ºC? (The specific heat of iron is 0.449 J/g·ºC) A. 4.77 J B. 4150 J C. 4210 J D. 29,500 J E. 33,700 J -9- Chapter 7 52. Which of the following types of electromagnetic radiation has the lowest frequency? A. red light B. green light C. blue light D. violet light E. ultraviolet radiation 53. Upon electrification, hydrogen produces the following line spectrum in the visible region of the electromagnetic spectrum. The light emitted for each line corresponds to transitions from the third (n=3), fourth (n=4), fifth (n=5), or sixth (n=6) energy level down to the second (n=2). Which color of light corresponds to the transition of an electron from the n=6 to the n=2 energy level? Color Wavelength (λ), nm Indigo 410.1 Blue 434 Blue-green 486 Red A. B. C. D. E. 656.3 indigo blue blue-green red all these colors are observed 54. Light with a wavelength of 650.0 nm appears as red light to the human eye? What is the energy of this light in joules? A. 1.291 × 10–13 B. 3.056 × 10–19 C. 3.056 × 10–28 D. 4.307 × 10–31 E. 4.612 × 1014 55. How many valence electrons does an atom of phosphorus have? A. 0 B. 3 C. 5 D. 8 E. 15 56. Which of the following elements requires the most energy to remove an electron from a neutral atom? A. Li B. Na C. K D. Rb E. Cs 57. Which of the following represents the electron configuration of cobalt? A. [Ar]3s23p9 B. [Ar]3s23d9 C. [Ar]3d7 D. [Ar]4s23d7 E. [Ar]4s24p9 58. How many orbitals are in a d subshell? A. 1 B. 3 C. 5 D. 6 E. 10 -10- 59. Which of the following orbital diagrams represents the correct ground state electron configuration for the oxide ion? A. 1s 2s 2p 1s 2s 2p B. C. 1s 2s 2p 1s 2s 2p 1s 2s 2p D. E. 60. An element was found to have two electrons in its valence s subshell and five electrons in its valence p subshell. To which group on the periodic table does this element belong? A. IIA (2) B. IIIB (3) C. VIIB (7) D. VA (15) E. VIIA (17) 61. Which of the following species has the electron configuration 1s22s22p63s23p64s23d104p65s24d105p6? A. Se B. Mo C. Te D. Xe E. Po 62. Which of the following best represents a p orbital? B A C D 63. Which of the following species has the most electrons? A. P3– B. Ar C. K+ D. Ca2+ E. All have the same number of electrons 64. Which species has its valence electrons closest to the nucleus? A. N B. P C. As D. Sb E. All have the same distance to the valence electrons -11- Chapter 8 65. Which of the following compounds is expected to have the lowest boiling point? A. KNO3 D. Mg(NO3)2 E. H2 B. CsF C. NaCl 66. Which of the following statements best describes a covalent bond in a compound? A. B. C. D. E. It is composed of only nonmetals. It is composed of a solid elemental metal. The compound is composed a metal and a nonmetal. It is held together by oppositely charged ions. There is no general rule to predict covalent bonds. 67. Which of the following species requires more than one electron dot symbol to describe the distribution of electrons? D. CH4 E. Br2 A. H2 B. CO32– C. H2O 68. When the electron dot symbol is complete for formaldehyde, CH2O, how many bonding electrons are present? (The connectivity of the atoms is shown below.) O C H A. 2 B. 3 C. 4 H D. 6 E. 8 69. What is the electron pair geometry and the molecular shape, respectively, for PH3? A. B. C. D. E. tetrahedral, trigonal planar tetrahedral, trigonal pyramidal tetrahedral, bent trigonal planar, trigonal planar trigonal pyramidal, trigonal pyramidal 70. Which of the following compounds is not an exception to the octet rule? (In each case the first element in the formula is central to all the other atoms.) A. PCl5 B. NO2 C. CF4 D. SF6 E. NO 71. Which of the following elements is least electronegative? A. fluorine B. carbon C. boron D. cesium E. lithium -12- 72. In which of the following compounds is the bonding polar covalent? A. CsF B. CO2 C. F2 D. NaCl E. O2 73. Which of the following molecules is composed of polar bonds but has a zero dipole moment (i.e. is nonpolar)? A. Cl2 B. SO2 C. CO2 D. NF3 E. PCl3 74. In the complete electron dot structure for NF3, which statement is not true? A. B. C. D. E. The F—N—F angle is about 109.5º. Nitrogen has six bonding and zero non-bonding electrons. There are 26 valence electrons. All atoms have satisfied valence shells. The molecule has tetrahedral electron pair geometry and trigonal pyramidal molecular shape. 75. In which of the following are partial negative and positive charges correctly placed? A. B. C. δ– F—Fδ+ F—Fδ– δ– S—Fδ+ δ+ D. E. δ– N—Oδ+ N—Oδ– δ+ 76. The electron pair geometry and molecular shape in SO2, respectively, are A. B. C. D. E. tetrahedral, bent tetrahedral, trigonal planar tetrahedral, tetrahedral trigonal planar, trigonal planar trigonal planar, bent Chapter 9 77. A balloon is filled with H2 gas to a volume of 750.0 mL. The pressure of the gas in the balloon is 781.2 mmHg, and the temperature is 23.7ºC. How many moles of H2 gas are in the balloon? A. 4.38 × 10-4 mol B. 6.50 × 10-2 mol C. 0.0316 mol D. 0.130 mol E. 0.396 mol 78. A balloon containing 20.0 L of air exerts a pressure of 2.00 atm. If the balloon is compressed to 15.0 L at a constant temperature of 25ºC, what is the final pressure? A. 0.375 atm B. 0.667 atm C. 1.50 atm D. 2.67 atm E. 16.7 atm 79. The images below represent a portion of a container that has gas atoms. Which image represents gases at the lowest pressure? A. B. C. D. -13- E. 80. Two 10-L balloons are filled with different gases, helium and carbon dioxide. Which of the following is the same for the two balloons? A. Density B. Mass C. Volume D. Number of moles of gas E. Both C and D Chapter 10 81. Which of the following best describes London-dispersion forces? A. They are forces of attraction only between polar molecules. B. They are covalent bonds between any kind of molecule. C. They are forces of attraction that result from the shifting distribution of electrons, even in molecules that are nonpolar. D. They are forces of attraction only between molecules that exhibit hydrogen bonding. E. They are present only in ionic compounds. 82. Which of the following compounds has the weakest intermolecular forces? A. CH3–O–CH3 B. CH3CH2–OH C. CH3CH2CH3 D. H2 E. CH2Cl2 83. The graph below shows boiling point data for compounds containing hydrogen and a halogen. Which of the following explains why HCl, HBr, and HI have lower boiling points even though they are heavier molecules? A. B. C. D. E. The covalent bonding between HF molecules makes it easier for the molecules to escape into the gas phase. HCl, HBr, and HI exhibit strong London dispersion forces so their boiling points are higher than HF. HF exhibits hydrogen bonding so it takes more energy for the molecules to escape into the gas phase. HCl, HBr, and HI exhibit strong hydrogen bonding so their boiling points are higher than HF. HF has the largest molar mass. 84. Which of the following compounds is expected to have the highest vapor pressure? (Hint: consider the graph in the previous question.) B. H2 B. HF C. HCl D. HBr E. HI 85. Which of the following exhibits dipole-dipole intermolecular forces? A. CH2F2 B. CO2 C. CBr4 D. CCl4 E. None of the above -14- 86. Which of the following compounds is most soluble in carbon tetrachloride, a nonpolar solvent? A. Br2 B. KBr C. CH3Cl D. KCl E. CH3CH2OH 87. Water vapor condenses when cooled below 100ºC? Gaseous ammonia condenses at –196ºC. What can we conclude about these two compounds based on these observations? A. B. C. D. E. Ammonia has stronger intermolecular forces than water. The hydrogen bonding in ammonia is stronger than the hydrogen bonding in water. The hydrogen bonding in ammonia is weaker than the hydrogen bonding in water. Water exhibits a higher vapor pressure than ammonia. Water exhibits hydrogen bonding but ammonia does not. Chapter 11 88. Which solution contains the highest concentration of ions? A. B. C. D. E. 1.0 M KCl 1.0 M K2CO3 1.0 M K3PO4 1.0 M NaCl All contain the same concentration of ions 89. What is the total concentration of ions in a solution that contains 0.250 M sodium sulfide, Na2S? A. 0.125 M B. 0.250 M C. 0.500 M D. 0.750 M E. 11.5 M 90. What is the molarity of a solution prepared by dissolving 235 g of KClO3 (MM = 122.55 g/mol) in enough water to give 3.5 L of solution? A. 0.0585 M B. 0.548 M C. 0.630 M D. 17.1 M E. 67.1 M 91. Which of the following will increase the solubility of a gas in solution? A. B. C. D. E. Increasing the gas pressure and increasing the temperature Increasing the gas pressure and decreasing the temperature Decreasing the gas pressure and decreasing the temperature Decreasing the gas pressure and increasing the temperature Adding more water 92. When CO2 dissolves in water, which of the following statements best describes the endothermic processes that occur? A. B. C. D. E. There is no energy gain or loss in the processes involved in CO2 dissolving. Energy is absorbed to break the ionic bonds in CO2. Energy is absorbed to break the ionic bonds in CO2 and water. Energy is absorbed to break the intermolecular forces in CO2 and the hydrogen bonds in water. Energy is absorbed when new intermolecular forces form between CO2 and water. -15- 93. How would you prepare 25.0 mL of 4.00 M NaCl solution for a stock solution that is 4.50 M? A. B. C. D. E. Add a little water to the 4.50 M solution. Measure 28.1 mL of 4.50 M solution and evaporate it until the volume reaches 25.0 mL. Mix 4.5 mL of 4.50 M solution with 20.5 mL of water. Measure 22.2 mL of 4.50 M solution and dilute it to a total volume of 25.0 mL. Measure 28.1 mL of 4.50 M solution and dilute it to a total volume of 25.0 mL. Chapter 12 The image below represents a hypothetical reaction X → Y. Use the image to answer the next two questions. C A B D 94. Which line corresponds to the energy difference between the reactants and products? 95. Which of the following statements best explains the reaction? A. B. C. D. E. The reaction is endothermic and can be described by the equation: X + Heat → Y The reaction is endothermic and can be described by the equation: X → Y + Heat The reaction is exothermic and can be described by the equation: X + Heat → Y The reaction is exothermic and can be described by the equation: X → Y + Heat Energy gain or loss cannot be predicted from the diagram. 96. Which of the following will affect the rate of a chemical reaction? A. B. C. D. E. The orientation of reacting species. The concentration of reacting species. The presence of a catalyst. The temperature. All of the above affect the rate of a reaction -16- 97. The equilibrium constant for the dissociation of an ionic compound is represented as Ksp, and corresponds to the equilibrium reaction that represents the dissolving process. For example, the equilibrium constant for the process FeS(s) ⇋ Fe2+(aq) + S2–(aq) is 6.0 × 10-19. The equilibrium constant for a reaction of this type indicates how soluble the compound may be. Given the Ksp values, which of the following compounds is most soluble? A. FeS (Ksp = 6.0 × 10-19) B. CuS (Ksp = 6.0 × 10-37) C. ZnS (Ksp = 3.0 × 10-23) D. PbS (Ksp = 3.4 × 10-28) E. HgS (Ksp = 4.0 × 10-54) 98. Ammonia can be prepared according to the reaction: N2(g) + 3H2(g) → 2NH3(g) Which of the following represents the equilibrium constant (Keq) expression for the reaction? A. K eq = [NH3 ] [N 2 ][H 2 ] D. K eq = [N 2 ][H 2 ] [NH 3 ] 2 B. K eq = C. K eq = [NH 3 ] 3 [N 2 ][H 2 ] E. K eq = [NH 3 ]2 ×[N 2 ][H 2 ]3 [N 2 ][H 2 ]3 [NH 3 ]2 99. At 350°C hydrogen gas reacts with gaseous iodine to form hydrogen iodide according to the reaction: H2(g) + I2(g) → 2HI(g) Keq = 60 (at 350°C) If at a particular point in time the concentrations of H2 and I2 are both 0.010 M and the concentration of HI is 0.10 M, which of the following statements best describes the system. A. B. C. D. E. [H2] = [I2] = [HI]2 so the reaction is at equilibrium. The reaction will shift to the right to reach equilibrium. The reaction will shift to the left to reach equilibrium. Keq >> 1 so reactants are always favored for this reaction. Keq << 1 so reactants are always favored for this reaction. 100. The following reaction is exothermic: 2SO2(g) + O2(g) → 2SO3(g) Which of the following will increase the amount of SO2 produced? A. Adding O2 B. Removing SO3 C. Increasing the temperature D. Decreasing the temperature E. Crossing your fingers -17- Sample Exam Answers Question Answer 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 C B B E A E E C B D A E A B A B E C D D C D A B D B E E B E B A C 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60 61 62 63 64 65 66 67 68 B E C C A C A E A D D B C C C C C D A A B C A D C E E D D E A E A B E -18- 69 70 71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 87 88 89 90 91 92 93 94 95 96 97 98 99 100 B C D B C B E E C D B E C D C A A A C C D B B D D D A E A B C C