An Anthology of General Chemistry Questions - Review Sheet | CHEM 111, Study notes of Chemistry

Download An Anthology of General Chemistry Questions - Review Sheet | CHEM 111 and more Study notes Chemistry in PDF only on Docsity! AN “ANTHOLOGY” OF GENERAL CHEMISTRY QUESTIONS This collection of questions may be of assistance to you in preparing for general chemistry examinations. Most of the questions provided are typical of ones used in multiple-choice general chemistry exams. Some of these questions may appear on the exams, modifications of some of these questions may appear on the exams, or questions similar to these may appear on the exams. In any case do not assume that only these questions will be used on exams. Some of the questions in the “Anthology” are designed not as exam questions but as aids to your learning of the material. The questions are grouped together by topic. The answers are provided in a separate section of the chemistry/biochemistry home page (“Anthology” Answers). Because the coverage of material differs from one general chemistry course to another, e.g. a topic covered in Chem 111 may not be covered in Chem 110, you should correlate your choice of topics for review in this “Anthology” with your lecture notes. These questions serve only to supplement not replace your lecture notes and your text. CONTENTS: Topic Question Numbers Scientific Method 1 - 3 Metric System and Units 4 - 12 Conversions and Conversion Factors 13 - 20 Scientific (Exponential) Notation 21 - 23 Significant Figures 24 - 33 Physical and Chemical Properties 34 - 35 Atomic Structure 36 - 38 Elements and Compounds 39 - 42 Molecules and Ionic Compounds 43 - 47 Polyatomic Ions 48 - 51 Naming Ionic Compounds 52 - 56 Naming Molecular Compounds 57 - 59 Balancing Chemical Equations 60 - 63 Electrolytes 64 - 68 Acids and Bases 69 - 73 Water-Insoluble Compounds 74 - 76 Ionic Equations 77 - 82 Reactions in Aqueous Solution 83 - 87 Other Reactions 88 - 89 The Mole Concept 90 - 91 Atomic Weights 92 - 95 Molecular and Formula Weights 96 - 104 Mass Percent 105 - 107 Empirical Formulas/Their Calculation 108 - 111 Molecular Formulas/Their Calculation 112 - 113 Elemental Analysis 114 - 115 Mole Concept/Chemical Equations 116 - 123 Stoichiometry(Chemical Arithmetic) 124 - 134 Stoichiometry - Solutions 135 - 141 Energy-Forms and Units 142 - 148 Enthalpy 149 - 153 Calorimetry 154 - 162 Thermochemical Equations 163 - 172 Hess Law 173 - 176 “Light”, Electromagnetic Radiation 177 - 185 Quantum Mechanics - General 186 - 189 Quantum Numbers of the H Atom 190 - 205 Many-Electron Atoms 206 - 208 Orbital Diagrams 209 - 210 Electron Configurations 211 - 222 9. The prefix used to indicate 106 is A. milli- B. mega- C. micro- D. mini- 10. The quantity 0.150 meter is specified by: A. 150 mm B. 150 cm C. 150 km D. 150 nm 11. The quantity 10 cm3 is identical to A. 10 L B. 10 mL C. 10 dL D. 10 kL 12. An object has a mass of 98.2 g. and occupies 71.7 mL. Its density equals A. 1.37 g/cm3 B. 7041 g.mL. C. 1.37 g/m3 D. 7071 g/mL CONVERSIONS AND CONVERSION FACTORS 13. A temperature of -40 degrees F equals A. -40 degrees C B. 233 K C. both A. and B. 14. The temperature in Las Cruces is reported to you as 41oC . You are more familiar with the Fahrenheit scale. Therefore, you calculate this temperature to be A. 72oF B. 106oF C. 89oF D. 97oF 15. An object has a volume of 0.0010 m3. Its volume given in cm3 is A. 0.10 B. 1000 C. 100 D. 10 16. A distance is determined to be 1000 µm. This distance can also be reported as A. 0.0010 m B. 0.100 cm C. 0.010 dm D. A,B, and C 17. Given that 2.54 cm = 1 in (exact) convert 75 miles to kilometers. A. 47 B. 92 C. 121 D. 109 18. A liquid has a density of 0.89 g/mL. Calculate the volume of 1.0 kg. of this liquid. A. 1.12 L B. 890 mL. C. 1.12 mL D. 890 L 19. Given that 454 g. = 1 lb., a 50.0 lb. weight can be labeled ---------- kg. A. 110 B. 31.8 C. 22.7 D. 89.5 20. You need to fill your gas tank in Mexico by purchasing 11.6 gallons of gas. You thus will purchase------- liters of gas. (1 qt.=.9464 L) A. 11.0 L B. 12.3 L C. 43.9 L D. 46.4 L SCIENTIFIC(EXPONENTIAL) NOTATION 21. Convert 12345 to scientific notation. A. 1,2345 B. 1.2345 C. 1.2345 x 105 D. 1.2345 x 104 22. Convert 12300 to scientific notation. A. 1,2300 B. 1.2300 C. 1.2300 x 104 D. 1.23 x 104 23. Convert 0.00010 to scientific notation. A. 1 x 10-3 B. 1.0 x 10-3 C. 1 x 10-4 D. 1.0 x 10- 4 SIGNIFICANT FIGURES 24. The number of significant figures in 12000 A. is 5 B. is 4 C. is 3 D. is 2 E. cannot be specified. 25. The number of significant figures in 0.001 A. is 4 B. is 3 C. is 2 D. is 1 E. cannot be specified. 26. The number of significant figures in 0.010 A. is 4 B. is 3 C. is 2 D. is 1 E. cannot be specified. 27. The number of significant figures in 1.01010 is A. 6 B. 5 C. 4 D. 3 E. 2 For questions 28 - 32 select the answer to the calculation that contains the correct units and the result to the correct number of significant figures. 28. The result of 38.7 cm x 9.1 cm is A.352 cm2 B. 352 cm C. 350 cm D. 350 cm2 29. The result of 30.70 kg x 1.3 m x 1.3 m/60.0 s x 60.0 s is A. 0.01 J B. 0.014J C. 0.0144 J D. 0.01441 J 30. The result of 0.9979 + 0.91 is A. 1.9079 B. 1.908 C. 1.91 D. 1 .9 31. The result of 1.009544 - 1.0093 is A. 0.0002 B. 0.00024 C. 0.000244 D. 0.0002440 32. The result of (3.6 x 104)(1.222 x 10-2) is A. 4.3992 x 10 2 B. 4.399 x 102 C. 4.40 x 102 D. 4.4 x 102 33. The result of (1.0 x 10-6 )(5.0 x 1011)/2.0 x 10-8 is A. 2.5 x 10 -74 B. 2.5 x 10 25 C. 2.5 x 10 13 D. 2.5 x 10 - 3 PHYSICAL AND CHEMICAL PROPERTIES 34. From the following statements, select that one or more that specifies a physical property. i. Sulfur burns in air. ii. Chlorine is a colored gas. iii. Ether evaporates at room temperature. iv. Iron rusts readily in the tropics. v. Water freezes at 0oC. vi. Mercury, a liquid, combines with oxygen , a gas, to produce a red solid. A. i.,ii.,and vi. B. ii., iii., and vi. C. ii.,iii.,v.,and vi. D. i.,iv.,and vi. 35. From the list given in Question 34, select one or more chemical properties. A. i., ii., and iii B. i., iv., and vi. C.i. and iv. D. i. and vi. ATOMIC STRUCTURE 36. An atom has an atomic number of 12 and a mass number of 24. This atom consists of A. 24 protons, 24 electrons, and 12 neutrons. B. 12 protons, 12 electrons, and 24 neutrons. C. 12 protons, 12 electrons, and 12 neutrons. 37. A given atom has an atomic number of 35. It is an atom of the element A. chlorine B. bromine C. boron D. calcium 55. The formula of copper(I) phosphate is A. Cu3PO4 B. Cu3(PO4)2 C. Cu(PO4)3 D. Cu2PO4 56. The correct name of TiO2 is A. titanium oxide B. titanium dioxide C. titanium(IV) oxide NAMING MOLECULAR COMPOUNDS 57. The numerical prefix indicating 4 is A. tri- B. tetra- C. tris- D. trig- 58. The common name of dihydrogen monoxide is A. peroxide B. ammonia C. water 59. Which of the following compounds is named incorrectly ? A. PCl3 “phosphorus trichloride” B. S2Cl2 “sulfur chloride” C. N2O “dinitrogen monoxide” D. P4S3 “tetraphosphorus trisulfide” BALANCING CHEMICAL EQUATIONS 60. When the following equation is balanced, the sum of all the balancing coefficients is: C3H8 + O2 → CO2 + H2O A. 13 B. 12 C. 11 D. 10 61. In balancing Mg3N2 + HCl → MgCl2 + NH3 the coefficient ---- must appear in front of HCl. A. 3 B. 6 C. 9 D. 12 62. When the following equation is balanced, the coefficient in front of the Na is Na + H2O → NaOH + H2 A. 1 B. 2 C. 3 D. 4 63. When the following equation is balanced, the sum of the balancing coefficients is: Al + Fe3O4 → Fe + Al2O3 A. 20 B. 22 C. 24 D. 26 ELECTROLYTES 64. Which general statement about electrolytes is correct ? A. An electrolyte is a substance that produces ions in aqueous solution. B. An electrolyte is a substance that in water produces a solution which conducts electricity. C. An electrolyte may be an ionic or molecular substance. D. All of the above. 65. Which of the following is a weak electrolyte ? A. HCl B. HCN C. NaOH D. HNO3 66. Which of the following is a strong electrolyte ? A. CH3NH2 B. HC2H3O2 C. HClO D. HClO4 67. Which of the following is a nonelectrolyte ? A. H2O B. NaCl C. table sugar D. Ca(OH)2 68. When NH3 dissolves in H2O, the resulting solution contains a limited number of ammonium and hydroxide ions but mostly ammonia molecules. This observation enables us to classify ammonia as a A. strong electrolyte B. weak electrolyte C. nonelectrolyte ACIDS AND BASES 69. According to the Bronsted-Lowry definition, an acid A. is an electron-pair donor B. is an electron-pair acceptor C. is a proton donor D. is a proton acceptor 70. Which one of the following species is not a strong acid ? A. HF B. HCl C. H2SO4 D. HNO3 71. HClO4 is classified as a strong acid because A. it is highly corrosive B. it reacts vigorously with H2O C. it ionizes essentially completely in aqueous solution. D. it reacts only with strong bases 72. Identify which of the following is a neutralization reaction. A. Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) B. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) C. CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) D. CaCO3(s) → CaO(s) + CO2(g) 73. Which one of the following is a strong base ? A. NH3 B. CH3NH2 C. NH2OH D. KOH WATER INSOLUBLE COMPOUNDS 74. Which one of the following statements about solubility is NOT true? A. All compounds containing Na+, NH4+, K+, or Rb+ are water soluble. B. All acetates and nitrates are water soluble. C. With a few exceptions, all sulfides, carbonates, and phosphates are insoluble in water. D. All chlorides, bromides, and iodides are water soluble. 75. Which one of the following is water soluble? A. Pb(C2H3O2)2 B. PbS C. PbSO4 D. PbCO3 76. To form a water-insoluble substance, a solution containing Ca2+ ions should be mixed with a solution containing ----- -- ions. A. Cl- B. SO42- C. NO3- D. ClO4- IONIC EQUATIONS 90. The mole, the chemistry counting unit, consists of A. 12 units B. 6.022 units C. 6.022 x 10-23 units D. 6.022 x 1023 units 91. The mole is defined in terms of the number of atoms of C-12 in exactly -------- grams of C-12. A. 1.0 B. 12.0 C. 100.0 D. 1000.0 ATOMIC WEIGHTS 92. The atomic masses provided on the Periodic Table A. are relative masses B. are based upon the C-12 mass as the reference mass C. are given in atomic mass units (amu) D. are based upon the proportion of each isotope present in the naturally occurring mixture. E. all of the above 93. A certain element has 2 isotopes, one having a mass of 84.9118 amu and a fractional abundance of 0.7215 and the other having a mass of 86.9092 and a fractional abundance of 0.2785. The atomic weight of this element is A. 85.9105 amu B. 85.4681 amu C. 86.0025 amu D. 85.7253 amu 94. The average atomic mass of B is 10.811 amu. It consists of two isotopes B-10 (10.013 amu) and B-11 (11.009 amu). Calculate the fractional abundance of B-10. A. 0.199 B. 0.801 C. 0.204 D. 0.796 95. The compound ammonia consists of 3 H atoms and 1 N atom per molecule, i.e. NH3. A sample of ammonia contains 7.933 g of N and 1.712 g of H. What is the atomic mass of N relative to that of H? A. 13.90 B. 14.01 C. 4.634 D. 13.96 MOLECULAR AND FORMULA WEIGHTS 96. The molecular weight of CO2 is A. 44.01 amu B. 28.01 amu C. 44.01 g D. 28.01 g 97. The molar mass of CO2 is A. 44.01 amu B. 28.01 amu C. 44.01 g D. 28.01 g 98. The molecular weight of C2H6O is A. 30.08 amu B. 46.08 amu C. 30.08 g D. 46.08 g 99. The molar mass of C2H6O is A. 30.08 amu B. 46.08 amu C.30.08 g D. 46.08 g 100. The molecular weight of N2 is A. 28.014 amu B. 28.014 g C.14.007 g D. 14.007 amu 101. The molecular weight of S8 is A. 32.066 amu B. 256.53 amu C. 32.066 g D. 256.53 g 102. Which of the following statements is incorrect ? A. The formula weight of S is 32.066 amu B. The molar mass of S is 32.066 g C. The formula weight of O2 is 32.0 g D. The molar mass of O is 32.0 g 103. The formula weight of NaNO3 is A. 85.0 amu B. 85.0 g C. 69.0 amu D. 69.0 g 104. The formula weight of Na2B4O7 is A. 185.2 amu B. 201.24 amu C. 174.44 amu D. 197.81 amu MASS PERCENT 105. The mass percent of oxygen in Na2B4O 7 is A. 22.86% B. 55.65% C. 61.52% D. 69.95% 106. A sample of Na2B4O7 contains 0.3478 g. of Na. What is the mass of this sample? A. 1.521 g B. 1.784 g C. 2.011 g D. 1.967 g 107. The percent composition of CH2O is -----%C, ----- %H, and -----%O. A. 37.2, 6.73, 56.1 B. 42.3, 6.51, 51.2 C. 39.8, 6.34, 53.9 D. 40.0, 6.73, 53.3 EMPIRICAL FORMULAS/THEIR CALCULATION 108. A compound contains only N and O. A sample of this compound contains 0.2801 g of N and 0.1600 g of O. What is the empirical formula of this compound? A. NO2 B. NO C. N2O D. N2O3 109. A compound has the percent composition 80.0% C and 20.0% H. Its empirical formula is A. CH B. CH3 C. CH2 D. CH4 110. Calculate the empirical formula of a compound with the composition 17.5% Na, 39.7% Cr, and 42.8% O. A. NaCrO4 B. Na2CrO4 C. Na2Cr2O7 D. NaCrO3 111. Calculate the empirical formula of a compound having the following % composition: C, 68.8%; H, 5.0%; O, 26.2%. A. C7H6O2 B. C4H3O C. C3H3O D. C4H3O2 MOLECULAR FORMULAS/THEIR CALCULATION 112. A compound is analyzed and found to have the empirical formula CH2 . Its molar mass is found to be 84 g/mole. Its molecular formula is A. C2H4 B. C4H8 C. C6H12 D. C8H16 113. A sample of a carbon-hydrogen-oxygen compound contains 1.69 mg of C, 0.285 mg of H, and 2.27 mg of O. Its molecular weight is 60.0 amu. Its molecular formula is A. CH2O B. C2H4O2 C. C3H8O 128. KO2 in the amount of 0.500 kg. is allowed to react with excess H2O. How many kg. of KOH are produced ? A. 0.500 B. 0.789 C. 0.635 D. 0.395 129. In this reaction 10.0 g of KO2 and 10.0 g of H2O are used. What mass in grams of KOH can be produced ? A. 7.91 B. 20.0 C. 6.22 D. 11.1 130. For the reaction as described in Ques. 129, one reactant was present in excess. How much of this reactant was left over ? A. 1.27 g B. 8.72 g C. 5.63 g D. 4.59 g 131. When 25.0 g of KO2 was allowed to react with excess H2O, 15.7 g of KOH was actually obtained. Calculate the percent yield. A. 100% B. 79.7% C. 15.7% D. insufficient information THE NEXT THREE QUESTIONS RELATE TO THE BALANCED EQUATION 3TiO2(s) + 4C(s) + 6Cl2(g) → 3TiCl4(g) + 2CO2(g) + 2CO(g) 132. TiO2 in the amount of 10.0 g is allowed to react with excess C and Cl2. The total mass of the products is A. 23.7 g B. 27.4 g C. 29.7 g D. cannot be determined 133. How much Cl2 must be provided to react with 10.0 g of TiO2 in the presence of excess C ? A. 17.7 g B. 8.86 g. C. 15.1 g. D. 18.6 g 134. How much of the excess C must react under conditions described in Question 133 ? A. 4.11g. B. 2.00 g. C. 13.3 g D. 7.17 g STOICHIOMETRY - SOLUTIONS 135. A solution is prepared by dissolving 10.0 g of NaOH in sufficient H2O to produce 250 mL of solution. Calculate the molarity of this solution. A. 0.250 M B. 0.500 M C. 0.750 M D. 1.00 M 136. A solution is prepared by dissolving 25.0 g. of ZnSO4 in 100.0 g. of H2O. This solution has a density of 1.124 g./mL. Calculate the molarity of this solution. A. 1.24 M B. 1.39 M C. 1.08 M. D. 1.67 M 137. A solution in the amount of 2.00 L is prepared by diluting 28.0 mL of 18.0 M H2SO4 with distilled H2O. Calculate the molarity of this solution. A. 0.252 M B. 0.504 M C. 0.126 M D. 0.630 M 138. For a certain reaction 0.050 mole of HCl is needed. What volume of 0.888 M HCl(aq) is needed ? A. 34.2 mL B. 17.9 mL C. 42.2 mL D. 56.3 mL. 139. What volume of 0.111 M AlCl3(aq) is needed to provide 0.010 mole of Cl- ? A. 30.0 mL. B. 60.0 mL C. 90.1 mL D. 120. mL 140. For the reaction HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l), how many mL of 0.0100 M NaOH are necessary to neutralize the HCl present in 10.0 mL of 0. 0900 M HCl ? A. 90.0 mL B. 9.00 mL C. 0.90 mL D. 900 mL 141. In an acid - base titration in which NaOH neutralized H2SO4: 2NaOH(aq) + H2SO4(aq) → 2H2O(l) + Na2SO4(aq), 30.0 mL of the acid of unknown concentration required 37.2 mL. of 0.1011 M NaOH for neutralization. Calculate the molarity of this sample of H2SO4(aq). A. 0.125 M B. 0.0627 M C. 0.0313 M D. 0.157 M ENERGY - FORMS AND UNITS 142. Which of the following statements about energy is correct ? A. Energy is neither created nor destroyed. B. Energy is either kinetic or potential. C. Energy readily changes from one form to another. D. All of the above. 143. When you start your car, A. electrical energy is converted to mechanical energy. B. chemical energy is converted to electrical energy. C. electrical energy is transferred as heat. D. chemical energy is transferred as heat. 144. The first law of thermodynamics states A. the energy of the universe is constantly increasing. B. the energy of the universe is a constant. C. the energy of the universe is constantly dissipating 145. The S.I. unit of energy is A. the calorie B. the erg C. the Joule D. the dyne 146. The Joule equals one A. g.m2/s2 B. g.cm2/s2 C. kg.m/s D. kg.m2/s2 147. Convert 5.1 cal to J. A. 1.22 kJ B. 1.22 J C. 21.3 kJ D. 21.3 J 148. An object has a mass of 150 g. It is moving at a speed of 9.95 m/s. Calculate its kinetic energy in J. A. 7430 J B. 7.43 J C. 14.9 J D. 14900 J ENTHALPY 149. Energy can be transferred as either heat or work. Enthalpy is the energy A. transferred as heat in a change occurring at constant volume. B. transferred as work in a change occurring at constant volume. C. transferred as heat in a change occurring at constant pressure. D. transferred as work in a change occurring at constant pressure. 150. The symbol for enthalpy is A. E B. U C. S D. H E. G 151. In an exothermic process the system A. absorbs energy and ∆H is positive B. absorbs energy and ∆H is negative C. releases energy and ∆H is positive D. releases energy and ∆H is negative. 152. Enthalpy, temperature, pressure are all state functions, but heat and work are not because FOR THE NEXT TWO QUESTIONS YOU WILL NEED TO USE THE STANDARD HEATS OF FORMATION GIVEN IN THE THERMODYNAMICS TABLE OF YOUR TEXT 167. What is the heat of reaction for P4(s) + 5O2(g) → P4O10(s) ? A. 2940J B. -2940J C. 2940kJ D. -2940kJ 168. Calculate the heat of reaction for: P4O10(s) + 6H2O(l) → 4H3PO4(aq) A. -113.2kJ B. -498kJ C. 113.2kJ D. 498kJ 169. In the reaction described by 2NO(g) + O2(g) → 2NO2(g), thermal energy in the amount of 114 kJ is released. Calculate ∆H per gram of NO. A. -1.9kJ B. 1.9kJ C. -3.8kJ D. 3.8kJ 170. Based upon the following thermochemical equation, calculate the grams of C3H 8(g) that must be burnt to provide 255kJ of thermal energy: C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g) ∆H = -2044kJ A. 4.95 g B. 5.50 g C. 3.78 g D. 9.11 g AGAIN USE TABULATED STANDARD HEATS OF FORMATION TO DO THE NEXT TWO QUESTIONS. 171. Which of the following liquids has the greatest molar heat of vaporization ? A. CCl4 B. CH3OH C. CS2 D. HCN 172. Using the information provided in this thermochemical equation and tabulated standard heats of formation, calculate the standard heat of formation of C4H10(g). C4H10(g) + 13/2 O2(g) → 4CO2(g) + 5H2O(l) ∆H = -2855 kJ A. 148 kJ B. -148 kJ C. -60.0 kJ D. 60.0 kJ HESS LAW 173. Using I H2O2(l) → H2O(l) + 1/2O2(g) ∆H=-98.0kJ and II 2H2(g) + O2(g) → 2H2O(l) ∆H=-571.6kJ determine ∆H for H2(g) + O2(g) → H2O2(l) A. -187.8 kJ B. 187.8 kJ C. -473.6 kJ D. 473.6 kJ 174. Given the following data: I 2SO2(g) + O2(g) → 2SO3(g) ∆H = -196.7 kJ II SO3(g) + H2O(l) → H2SO4(l) ∆H = -130.1 kJ what is the heat of reaction for 2SO2(g) + O2(g) + 2H2O(l) → 2H2SO4(l) ? A. -456.9 kJ B. 456.9 kJ C. 326.7 kJ D. -326.7 kJ 175. Determine the heat of reaction for C2H5OH(l) + O2(g) → CH3COOH(l) + H2O(l) by using I C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(l) ∆H = -1366.9kJ II CH3COOH(l) + 2O2(g) → 2CO2(g) + 2H2O(l) ∆H = -869.9kJ A. -2237 kJ B. 2237 kJ C. -497 kJ D. 497 kJ 176. Given the following data: I N2(g) + 3H2(g) → 2NH3(g) ∆H = -91.8 kJ II C(s) + 2H2(g) → CH4(g) ∆H = -74.9 kJ III H2(g) + 2C(s) + N2(g) → 2HCN(g) ∆H = 270.3 kJ determine the heat of reaction for: CH4(g) + NH3(g) → HCN(g) + 3H2(g) A. -256 kJ B. 256 kJ C. 302 kJ D. -302 kJ “LIGHT”, ELECTROMAGNETIC RADIATION 177. The spectrum of light does not include A. X-rays B. gamma rays C. sound waves D. microwaves 178. Which of the following wavelengths corresponds to the most energetic light ? A. 750 nm B. 12.5 nm C. 7000 nm D. 27 nm 179. Which of the following equations specifies the energy of light ? A. E = λν B. E = hν C. E = 1/2mv2 D. E = mc2 180. Calculate the wavelength of light having a frequency of 5.00 x 1014 Hz. A. 600 nm B. 150 nm C. 60000 nm D. 15000 nm 181. Calculate the frequency of light having a wavelength equal to 626 nm. A. 1.88 x 104 Hz B. 4.79 x 1016 Hz C. 4.79 x 1014 Hz D. 1.88 x 107 Hz 182. The quantum of light energy is the A. neuron B. exciton C. photon D. phonon 183. A photon has a frequency of 4.20 x 1013 Hz. Calculate its energy in J. A. 1.58 x 10-21 J B. 2.78 x 10-20 J C. 6.34 x 10-20 J D. 3.60 x 10-9 J 184. A photon has a wavelength of 155 nm. Calculate is energy in J. A. 7.80 x 10-17 J B. 3.51 x 10-14 J C. 1.28 x 10 -18 J D. 4.27 x 10-11 J 185. What observations or series of experiments directly led to Einstein’s quantization of the energy of light ? A. the line spectra of atoms B. the photoelectric effect C. black-body radiation D. the scattering of alpha particles. QUANTUM MECHANICS - GENERAL 186. Which of the following statements is FALSE? A. Energy flows between objects in discrete indivisible amounts rather than continuously. B. The energy of the atom is restricted to specific amounts. C. Once the energy of the electron is known so is its exact position. D. Light can exhibit either particle or wave behavior depending upon the system with which it is interacting. 187. The concept that the energy of the atom is quantized was based upon a consideration of A. black-body radiation B. the line spectra exhibited by atoms C. the stability of the Rutherford (nuclear) atom D. A, B, and C 206. The Pauli exclusion principle states that in the many- electron atom A. the exact position of any one of the electrons cannot be given. B. no two electrons can have the same value of the spin quantum number . C. each electron must have its own unique set of the 4 quantum numbers . D. all of the above. 207. The maximum number of electrons that can be accommodated in the 4d subshell is A. 4 B. 6 C. 8 D. 10 208. The maximum number of electrons that can be accommodated in the n = 3 shell is A. 18 B. 16 C. 14 D. 12 ORBITAL DIAGRAMS 209. Which of the following abbreviated orbital diagrams is invalid ? s p d A. ↑↓ ↑↓ ↑↓ ↑↓ B. ↑↓ ↑↓ ↓↑ ↑↓ ↑ ↑ ↑ ↑ ↑ C. ↑↓ ↑↑ ↓↓ ↑↑ D. ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 210. Which of the following orbital diagrams DOES NOT represent the outermost electrons of an atom in its ground (lowest energy) state? s p d A. ↑↓ ↑ ↑↓ ↑↓ ↑↓ ↑↓ B. ↑↓ ↑ ↑ ↑ C. ↑↓ ↑↓ ↑↓ ↑↓ D. ↑↓ ↑ ↓ ↓ ELECTRON CONFIGURATIONS 211. Which of the following electron configurations is VALID ? A. 1s22s22p73s1 B. 1s22s22p62d2 C. 1s22s22p63s13p6 D. 1s12s12p8 212. Which of the following electron configurations represents the ground state of O ? A. 1s22p6 B. 1s22s22p23s2 C. 1s22s22p22d2 D. 1s22s22p4 213. Which of the following represents the ground state of potassium ? A. 1s22s22p63s23p63d1 B. 1s22s22p63s23p64s1 C. 1s22s22p62d103s1 D. 1s22s22p43s23p43d5 214. Which of the following represents the ground state of vanadium ? A. 1s22s22p63s23p63d34s2 B. 1s22s22p63s23p63d5 C. 1s22s22p43s23p43d9 D. 1s22s22p43s23p43d74s2 215. Which of the following represents the ground state of gallium ? A. [Ar]4s24p1 B. [Ar]3d104s24p1 C. [Ar]4s24p64d5 D. [Ar]4f13 216. Which of the following represents the ground state of Sm (samarium) ? A. [Xe]4f66s2 B. [Xe]5d66s2 C. [Xe]4f35d36s2 D. [Xe]4f35d5 217. Which of the following represents the ground state of silver ? A. [Kr]4d104f1 B. [Kr]4d94f2 C. [Kr]4d95s2 D. [Kr]4d105s1 218. Which of the following represents the ground state of the manganese(II) ion ? A. [Ar]3d5 B. [Ar]3d34s2 C. [Ar]3d44s1 D.[Ar]3d54s1 219. Which of the following represents the ground state of the oxide ion ? A. [Ne] B. 1s22s22p4 C. 2s22p6 D. [Ar] 220. From inspection of their of their orbital diagrams determine which of the following species is paramagnetic. A. Mg B. N3- C. Mn D. Na+ 221. The valence electron configuration of an atom is ns2. This atom is A. an alkali metal B. a halogen C. a rare earth D. an alkaline earth metal. 222. With one exception, a noble gas has the outermost electron configuration A. ns2np4 B. ns2np6 C. (n-1)d10ns2np6 D. (n+1)s2np6 ATOMIC AND IONIC RADII 223. The radius of an atom depends upon A. n, the principle quantum number, and Z, the atomic number B. l, the angular momentum quantum number, and n C. n and A, the mass number D. l and Z 224. Which of the following Mg, Ca, Ba, or Sr has the largest atomic radius ? A. Mg B. Ca C. Ba D. Sr 225. Which of the following has the smallest atomic radius ? A. Be B. N C. B D. O 226. Which of the following arrangements shows the atoms increasing from smallest to largest ? A. Cs<Rb<Mg<Si<C B. Rb<Cs<Si<Mg<C C. C<Si<Mg<Rb<Cs D C<Mg<Rb<Si<Cs 227. Which of the following is the largest species ? A. F- B. Cl- C. O2- D. S2- E. N3- F. P3- 241. Ionic compounds A. have high melting points. B. are always solids at normal conditions. C. conduct electricity only when melted or dissolved in water. D. all of the above. 242. Which of the following equations describes the lattice energy of a substance ? A. M(s) + energy → M+(g) + e- B. MX(s) + energy → M+(g) + X-(g) C. X-(g) + energy → X(g) + e- D. M+(g) + X-(g) → MX(g) + energy 243. When a representative (main-group) element forms an ion, it attains the ------- electronic configuration A. alkali metal B. noble gas C. pseudo-noble gas D. either B or C 244. Cations formed by representative elements have A. a positive charge equal to the group number. B. a charge equal to the group number - 8. C. one of several different positive charges. D. a negative charge equal to the group number. LEWIS DOT STRUCTURES 245. The correct Lewis dot structure of the carbon atom is: C C C CA. B. C. D. 246. The correct Lewis dot structure for N3- is: A. C.B.N N N NeD. 3- 3- 3- LEWIS STRUCTURAL FORMULAS 247. The formula of AlF3 in terms of Lewis dot structures is: D.C.Al Al F 33 FAl AlFA. F B. 248. A clue to the number of bonds an atom forms is provided by: A. the atom's position in the periodic table B. the number of electrons appearing singly about the atom's Lewis dot structure C. the atom's electronegativity D. the atom's electron affinity 249. The valid Lewis formula for NH3 is: A. B. C. D.HNH H HNH H HNH H HNH H 250. The correct Lewis formula for PCl3 is: ClClP ClPCl Cl ClPCl Cl Cl ClPCl Cl A. B. C. D. 251. The correct Lewis formula for F2NNF2 is: FNNF F F FNNF F F FNNF F F FNNF F F A. B. C. D. 252. The correct Lewis formula for H3COCH3 is: B. C. D. A. H C O C H H H H H H C O C H H H H H H C O C H H H H H H C O C HHH H H 253. A correct Lewis formula for NNO is: N N OA. B. D. A and B E. A and C N N O C. N N O F. B and C 254: The correct Lewis formula for CO2 is: OCO OCO OCOA. B. C. D.OCO 255. Which of the following is a valid Lewis structure for CO? OC OCA. B. C. D.OC OC 256. Which of the following cannot exhibit resonance? A. SO2 B. O3C. CH2O D. NO2- 257. Which of the following is a valid Lewis structure for HN3? 258. A student offers the following Lewis structure for HCN: H=C=N. What errors do you observe in the proposed structure ? A. a miscount in the number of valence electrons B. a doubly bonded hydrogen C. nitrogen without an octet D. all of the above 259. At least two valid Lewis structures can be drawn for N2O . 265. Which of the following molecules is “electron deficient” ? A. CH4 B. BH3 C. NH3 D. H2O 266. Which of the following molecules violates the octet rule ? A. CH4 B. CCl4 C. SF4 D. XeF4 E. C and D COVALENT BONDS 267. Covalent bonds form when A. metals react with nonmetals. B. metals react with one another. C. nonmetals including hydrogen react with one another. 268. Which statement about covalent bonding is correct ? A. The bond involves the sharing of one or more electron pairs. B. In the formation of the bond each of the bonded atoms can contribute a single electron. C. In the formation of the bond one of the bonded atoms can contribute both electrons. D. The electron pair or pairs can be shared equally by the bonded atoms. E. The electron pair or pairs can be shared unequally by the bonded atoms. F. All of the above. 269. In the formation of a covalent bond A. energy is absorbed. B. energy is released. C. energy is either absorbed or released depending upon the identity of the bonding atoms. 270. In the formation of the ammonium ion the ammonia molecule donates a pair of electrons to a hydrogen ion to form a covalent bond. The bond is b e s t described as A. a dating bond B. a polar bond C. a coordinate covalent bond. D. a nonpolar bond. ELECTRONEGATIVITY 271. The quantity termed “electronegativity” is a quantity A. artificially constructed. B. very easy to measure experimentally. C. very difficult to measure experimentally. D. difficult or easy to measure experimentally. Use the following data to answer the next three questions. Atom Electronegativity H 2.1 Cs 0.7 O 3.5 N 3.0 P 2.1 F 4.0 Cl 3.0 272. Predict which of the following bonds is the most polar. A. H-O B. H-P C. H-Cl D. H-F 273. Predict which of the following is ionic rather than covalent in nature. A. P-N B. H-Cs C. F-Cl D. N-F 274. Which of the following bonds is (are) nonpolar ? A. F-O B. O-O C. P-H D. A and B E. B and C BOND LENGTH, STRENGTH, ORDER 275. Bond order refers to A. the arrangement of the bonds within a molecule in order of their polarity. B. the arrangement of the bonds within a molecule in order of their length. C. the arrangement of the bonds within a molecule in order of their strength. D. the number of electron pairs shared by the bonded atoms. 276. In a series of compounds two nitrogen atoms are bonded to one another. In Compound A the N to N bond order is 3, in Compound B it is 2, and in Compound C, it is 1. A. The strongest N to N bond is in Compound C. B. The longest N to N bond is in Compound C. C. The weakest N to N bond is in Compound A. D. The N to N bond in Compound B is easier to break than the one in Compound C and harder to break than the one in Compound A. BOND ENERGY 277. For which of the following equations does ∆H represent a bond energy ? A. C(s) → C(g) B. NaCl(s) → Na+(g) + Cl-(g) C. HCN(g) → H(g) + CN(g) D. Br2(l) → 2Br(g) 278. From the following thermochemical equation determine the average C-H bond energy: C(g) + 4H(g) → CH4(g) ∆H = -1662 kJ A. 1662 kJ B. 416 kJ C. -1662 kJ D. -416 kJ The following data may be of use in answering the next two questions. BOND BOND ENERGY (kJ/mol) H - H 432 Cl - Cl 240 H - Cl 428 C - Cl 327 C - C 346 C = C 602 279. Estimate ∆H for CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g) A. -99 kJ B. +99 kJ C. -1411 kJ D. +1411 kJ 280. Estimate ∆H for H2C=CH2(g) + H2(g) → H3C-CH3(g) A. -191 kJ B. +202 kJ C. +272 kJ D. -144 kJ MOLECULAR GEOMETRY 281. The geometry of a molecule, i.e. its general shape, is determined by C. a square pyramid D. the letter “T” 298. According to VSEPR, ClF3 possesses a ------ shape. A. trigonal planar B. trigonal pyramidal C. “T” D. tetrahedral 299. According to VSEPR, CO2 is a -------- molecule. A. linear B. bent C. trigonal D. tetrahedral 300. According to VSEPR, H2CO possesses A. bent geometry B. trigonal planar geometry C. distorted trigonal planar geometry D. distorted tetrahedral geometry 301. According to VSEPR, H2CO possesses A. bent geometry B. trigonal planar geometry C. distorted trigonal planar geometry D. distorted tetrahedral geometry POLAR MOLECULES 302. Which of the following statements is incorrect ? A. An electric dipole consists of two partial electrical charges equal in magnitude but opposite in sign and separated by a distance. B. Whether a molecule is polar or not can be determined experimentally. C. A polar molecule must contain polar bonds. D. A polar molecule cannot possess a completely symmetric geometry. E. None of the above. 303. Which of the following geometries is not symmetric ? A. trigonal planar B. tetrahedral C. octahedral D. trigonal pyramid 304. Which of the following molecules is polar ? A. BeCl2 B. CO2 C. SO2 D. I3- 305. Which of the following molecules is polar ? A. SiH4 B. TeCl4 C. XeF4 D. AlF4- HYBRID ATOMIC ORBITALS ON CARBON 306. Hybrid atomic orbitals were “invented” because A. the hydrogen-like atomic orbitals were too few in number. B. the hydrogen-like atomic orbitals could not produce the correct molecular geometry. C. the energies of the hydrogen-like atomic orbitals did not match the energies of the molecules. 307. Which statement about the formation of hybrid atomic orbitals is incorrect ? A. Only orbitals close in energy combine with one another. B. The number of hybrid orbitals formed always equals the number of hydrogen-like atomic orbitals used. C. The hybrid orbitals formed have well-defined directional characteristics. D. None of the above. 308. The hybrid orbitals necessary to tetrahedral geometry about the carbon atom are termed A. sp hybrids B. sp2 hybrids C. sp3 hybrids D. sp4 hybrids 309. The hybrid orbitals on a doubly bonded carbon are A. sp hybrids B. sp2 hybrids C. sp3 hybrids D. sp4 hybrids. 310. Which set of hybrid orbitals meets the following description: two in number, at an angle of 180o from one another, used in the formation of a multiple bond between carbons. A. sp4 B. sp3 C. sp2 D sp 311. A carbon atom forms a triple bond by the use of A. two sp hybridized atomic orbitals. B. two sp hybridized atomic orbitals and one 2p hydrogen-like atomic orbital. C. one sp hybridized atomic orbital and two 2p hydrogen-like atomic orbitals. D. three sp2 hybridized atomic orbitals. 312. When the geometry about a carbon atom is trigonal planar, its four bonds consist of A. two sigma(σ) and two pi(π) bonds. B. four sigma bonds. C. three sigma and one pi bond. D. four pi bonds. THE GASEOUS STATE 313. A gas is A. a fluid B. highly compressible C. highly disordered D. all of these 314. Which of the following statements is correct ? A. The volume of any gas is the volume of its container. B. A gas exerts pressure due to collisions of the gaseous molecules against a surface. C. Under ordinary conditions, all true gases behave the same. D. All of the above. 315. Which of the following is not a unit of pressure ? A. atmosphere B. millimeters of Hg C. Torr D. manometers of Hg E. psi 316. The device commonly used to measure atmospheric pressure is A. a manometer B. a sphygmomanometer C. a torricelli D. a baromete r 317. Convert 151 mmHg to Torr. A. 0.199 Torr B. 15.1 Torr C. 151 Torr D. 199 Torr 318. A gas exerts a pressure of 1.78 atm. This pressure in mmHg is A. 1350 B. 1070 C. 0.556 D. 1.78 319. Mercury is commonly used in barometers because of its high density (13.596 g/cm3) . An alcohol having a density of 0.899 g/cm3 is substituted for mercury. Calculate the height of an alcohol column exerting the same pressure as a mercury column 15.6 mm high. A. 236 mm B. 189 mm C. 212 mm D. 273 mm EMPIRICAL GAS LAWS 335. A sample of a gas occupies 129 mL at 28oC. and exerts a pressure of 682 mmHg. How many moles of gas are present? A. 5.04 x 10-2 B. 4.68 x 10-3 C. 3.56 D. 6.31 x 10 - 3 336. A sample of N2(g) was generated in the lab. At 25oC and 692 Torr its volume was 97.3 mL. How many grams of N2 were produced ? A. 3.62 x 10-3 B. 1.01 x 10-1 C. 1.59 x 10-3 D. 3.99 x 10- 1 337. Oxygen gas is prepared in the laboratory via 2KClO3(s) → 2KCl(s) + 3O2(g). KClO3 in the amount of 10.0 g completely decomposed in this manner. What volume of O2(g) at 31oC and 0.897 atm was produced ? A. 3.40 L B. 2.26 L C. 0.346 L D. 1.55 L 338. For a given reaction 300.0 g of N2(g) is required. What volume of N2(g) at 25oC and 1 atm must be pumped into a reactor to provide the needed N2(g) ? A. 22.0 L B. 49.7 L C. 262 L D. 193 L 339. For a given reaction 300.0 g of N2(g) is required. What volume of N2(g) at 25oC and 1 atm must be pumped into a reactor to provide the needed N2(g) ? A. 22.0 L B. 49.7 L C. 262 L D. 193 L 340. A given number of moles of any gas, regardless of its identity, occupies a given volume at a specific combination of temperature and pressure. Calculate the number of moles of a gas present in 1 L at 25oC and 709 mmHg. A. 0.455 B. 0.0381 C. 0.0345 D. 0.596 341. Calculate the density of Cl2(g) at 25oC and 709 Torr. A. 2.70g/L B. 32.3g/L C. 2.45g/L D. 1.95g/L 342. Calculate the density of Cl2(g) at 37oC. and 0.592 atm. A. 2.63g/L B. 1.65g/L C. 2.70g/L D. 1.92g/L 343. Calculate the density of O3(g) at 25oC and 709 Torr. A. 2.70g/L B. 1.94g/L C. 1.83g/L D. 2.49g/L 344. Two gases at the same temperature and pressure can differ in their density because A. the sample of one gas can contain a greater number of molecules than does the sample of the other gas. B. the strength of the attractive forces between molecules are different in the two gases. C. the molecules of the two gases differ in their molecular volumes. D. the two gases differ in their molecular weight. GAS MIXTURES 345. In any mixture of nonreacting gases, A. the volume of each gas is the volume of the container. B. each gas is exerting its own pressure. C. the total pressure is the sum of each gas’ s individual pressure. D. the pressure each gas is exerting depends directly on the number of its molecules present in the mixture. E. all of the above. The next three questions relate to a mixture of noble gases consisting of 0.150 mole of He, 0.450 mole of Ne, and 0.300 mole of Ar. 346. The mole fraction of Ar in this mixture is A. 0.300 B. 0.333 C. 0.667 D. 0.500 347. The partial pressure of the Ne in the mixture is 275 Torr. The total pressure of the mixture is A. 550 Torr B. 275 Torr C. 611 Torr D. 900 Torr 348. The partial pressure of the He is A. 183 Torr B. 275 Torr C. 91.7 Torr D. 157 Torr 349. The O2(g) generated by the decomposition of KClO3 was collected by the displacement of water at 31oC (vapor pressure of H2O = 32.0 mmHg). The “wet gas” had a volume of 75.8 mL . The barometric pressure during the collection process was 683mmHg. What pressure did the O2(g) exert ? A. 715 mmHg B. 683 mmHg C. 651 mmHg D. 667 mmHg 350. How many grams of O2(g) were generated in the process described in the previous question ? A. 2.60 x 10-3 B. 8.33 x 10-2 C. 2.55 x 10-2 D. 6.77 x 10- 3 KINETIC MOLECULAR THEORY OF GASES 351. The kinetic molecular theory of gases was developed to explain the behavior of A. real gases B. noble gases C. ideal gases D. inert gases 352. Which of the following statements is not a postulate of the kinetic molecular theory of gases. A. The volume of the molecules making up the gas is negligible compared to the distances between the molecules. B. The gaseous molecules move constantly and randomly in straight lines throughout the volume of the container. C. The attractive forces between molecules are essentially negligible. D. The average kinetic energy of the molecules is directly proportional to the pressure they exert. 353. The gaseous state is the simplest state of matter to study because the identity of the gas is irrelevant under ordinary conditions. This irrelevance is based upon A. the fact that the gaseous molecules are electrically neutral. B. the fact that molecular volume is immaterial under ordinary conditions. C. the fact that the gaseous molecules neither attract nor repel one another under ordinary conditions. D. A, B, and C E. B and C COMPARISON OF GASES, LIQUIDS, AND SOLIDS 354. The least compressible state of matter is the ---- state. A. liquid B. solid C. gaseous D. irrepressible 355. This substance is a fluid of relatively high density. It conforms to the shape but not the volume of its container. It is a A. gas B. solid C. liquid D. plasma 367. In a phase diagram, a line segment separates the liquid from the solid state. A point along this line designates a pressure of 6 atm and a temperature of -56oC. These data tell us that A. the substance is a pure solid at this T and P. B. the substance is a pure liquid at this T and P. C. at 6 atm the substance melts at -56oC D. at 6 atm the substance freezes at -56oC E. C and D. 368. The triple point of water is 0.01oC and 4.58 Torr. This information tells us that at a pressure less than 4.58 Torr, A. ice cannot exist B. steam cannot exist C. liquid water cannot exist PROPERTIES OF THE LIQUID STATE 369. A liquid’s resistance to flow is termed its A. surface tension B. viscosity C. cohesion D. capillarity 370. The viscosity of a liquid is -------- and its vapor pressure is ---------- by an increase in the temperature. A. increased, increased B. decreased, decreased C. increased, decreased D. decreased, increased 371. Although a needle can be made to float upon water, the addition of soap causes the needle to sink. This effect occurs because A. the soap lowers the viscosity of the water. B. the soap lowers the vapor pressure of the water. C. the soap lowers the surface tension of the water. D. the soap lowers the meniscus of the water. INTERMOLECULAR FORCES 372. Intermolecular forces A. are electrical in nature. B. are weaker than chemical bonds. C. vary in strength depending upon the distance between molecules. D. vary depending upon the nature and identity of the molecules. E. account for the formation of the liquid and solid states. F. all of the above 373. The intermolecular attractive forces in Substance I are stronger than those in Substance II. Therefore, A. I has a lower vapor pressure than does II B. I has a lower heat of vaporization than does II C. I has a lower melting point than does II D. I has a lower boiling point than does II E. all of the above 374. Identify the intermolecular force from its description: acts on all molecules, depends upon the occurrence of an instantaneous (“snapshot”) dipole which induces dipoles in surrounding molecules. A. dipole-dip ole interactions B. covalent bonding C. hydrogen bonding D. London (dispersion) forces 375. Only molecules containing a N, O, or F atom experience A. dipole-dipole interactions B. covalent bonding C. hydrogen bonding D. London(dispersion) forces 376. This type of intermolecular attraction has been described as an extreme form of dipole-dipole interaction. A. covalent bonding B. hydrogen bonding C. London forces D. ionic bonds 377. Which of the following species experiences only London forces ? A. HCl B. H2O C. NH3 D. CH3CH3 378. Which of the following is subject to dipole-dipole interactions? A. H2 B. He C. CO2 D. H2S 379. Consider an ammonia molecule. It experiences A. dipole-dipole interactions B. London forces C. hydrogen bonding D. all of the above 380. Which of the following experiences dipole-dipole interactions ? (HINT: review VSEPR) A. CH4 B. SiF4 C. TeCl4 D. XeF4 381. In general the strongest type of intermolecular force is A. dipole-dipole interactions B. covalent bonding C. hydrogen bonding D. London or dispersion forces 382. Consider the two molecules CH3-O-CH3, an ether, and the alcohol CH3CH2OH. A. The alcohol has a higher vapor pressure than does the ether. B. The alcohol has a lower freezing point than does the ether. C. The alcohol has a higher boiling point than does the ether. D. All of the above 383. Following the usual trend the boiling points of the hydrogen halides should increase systematically from the lowest for HF to the highest for HI. Instead, HF has the highest boiling point. Which of the these explanations for HF’s anomalous behavior is the accepted explanation? A. The HF molecule is so light in mass it possesses exceptionally high kinetic energy. B. The lone pairs on the F atom are drawn so close to the nucleus that the molecule experiences very weak attractive forces. C. The presence of the fluorine atom which is very highly polarizable causes the molecule to experience very strong London forces. D. The hydrogen atom of one HF molecule is strongly attracted by a lone pair of electrons on the fluorine atom of an adjacent HF molecule. TYPES OF SOLIDS 384. Amorphous solids differ from crystalline solids in that A. crystals have a well-defined orderly structure, amorphous solids do not. B. crystalline solids have sharp melting points, amorphous solids do not but soften over a temperature range. 398. Benzene is a liquid hydrocarbon. Predict which of the following is most likely to be soluble in benzene ? A. H2O(l) B. C6H14(l) C. CH3COOH(l) D. CH3OH(l) 399. When an ionic compound dissolves in water, the ----- is released. A.lattice energy B. energy of hydration C. vaporization energy D. bond energy 400. The two energy changes that influence the formation of aqueous solutions of ionic compounds are A. ionization energy and electron affinity B. ionization energy and hydration energy C. electron affinity and hydration energy D. electron affinity and lattice energy E. hydration energy and lattice energy 401. Which of the following cations experiences the greatest energy of hydration ? A. Li+ B. Na+ C. K+ D. Rb+ 402. Which of the following ions experiences the greatest energy of hydration ? A. Na+ B. Mg+2 C. Al3+ D. N3- 403. Which of the following compounds possesses the greatest lattice energy ? A. Mg(OH)2 B. Ca(OH)2 C. Sr(OH)2 D. Ba(OH)2 EFFECT OF T AND P ON SOLUTIONS 404. A gas is most soluble in a liquid at A. low temperature and low pressure B. high temperature and high pressure C. low temperature and high pressure D. high temperature and low pressure 405. High temperatures increase the solubility in water of most A. gases B. ionic compounds C. hydrocarbons 406. When ammonium nitrate dissolves in water the solution cools because the solution process is A. endoentropic B. exoentropic C. endothermic D. exothermic 407. Acetylene gas at room temperature dissolves to the extent of 320 g per liter of liquid acetone at 12 atm pressure. Determine the Henry’s law constant for acetylene in acetone at room tempera ture . A.0.0375 L atm/g B. 26.7 g/L atm C. 3840 g atm/L D. 0.260 L/g atm 408. Calculate the solubility of acetylene in acetone at room temperature and a pressure of 5 atm. A. 187 g/L B. 195 g/L C. 134 g/L D. 173 g/L SOLUTIONS - CONCENTRATION UNITS 409. Which of the following concentration units changes with changes in temperature ? A. mass percent B. mole fraction C. molarity D. molality E. ppm 410. A solution is prepared by dissolving 0.100 mole of HCl in 75.0 g of water. Calculate the mass percent HCl in this solution. A. 0.133% B. 4.64% C. 4.87% D. 4.01% 411. Any mixture of gases is a solution. Such a solution is prepared by mixing 10.0 g of He, 10.0 g of Ne, and 10.0 g of Ar. The mole fraction of Ar in this mixture is A. 0.0769 B. 0.333 C. 0.156 D. 0.00321 412. A solution is prepared by dissolving 10.0 g of HCl in sufficient water to produce 150 mL of solution. Calculate the molarity of this solution. A. 0.194M B. 1.83M C. 6.67M D. 3.22M 413. A solution is prepared by dissolving 20.0 g of NaOH in 750 g. of water. The molality of this solution is A. 26.7m B. 0.0267m C. 0.667m D. 0.000667m 414. A 50.0% by mass solution of HNO3(aq) has a density of 1.31g/mL. Calculate the molarity of this solution. A. 10.4M B. 14.7M C. 11.5M D. 9.77M 415. Calculate the molality of the 50.0% nitric acid solution. A. 19.1m B. 16.7m C. 15.9m D. 10.4m 416. Calculate the mole fraction of nitric acid in 50.0% nitric acid solution. A. 0.794 B. 0.195 C. 0.500 D. 0.222 417. The antiseptic H2O 2(aq) is 0.655 molal in the hydrogen peroxide. Calculate the mole fraction of peroxide present. A. 0.0118 B. 0.0116 C. 0.119 D. 0.211 COLLIGATIVE PROPERTIES 418. Colligative properties are the properties of a solution that depend upon A. the identity of the solute particles present in the solution. B. the size of the solute particles. C. the identity of the solvent of the solution. D. the size of the solvent particles. E. the number of solute particles present. F. the number of solvent particles present. 419. Colligative properties include A. vapor pressure lowering B. boiling point elevation C. freezing point depression D. osmotic pressure E. all of the above 420. A nonvolatile, nonelectrolyte in the amount of 0.0900 mole is dissolved in 100.0 g of water. Given that to 25oC, the vapor pressure of pure water is 23.8 Torr, calculate the vapor pressure of the solution. A. 23.8 Torr B. 0.379 Torr C. 23.4 Torr D. 0.369 Torr 421. Calculate the freezing point in oC of a solution containing 0.0100 mole of a nonelectrolyte in 100.0 g of water. A. -0.186 B. +0.186 C. 0.010 D. -0.010 422. Calculate the boiling point of the solution of Question 421.