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Chemistry 400: General Chemistry Discussion: Review for Multiple Choice Section of Final Chapter 3 1) Which of the following contains BOTH ionic and covalent bonds? A) CaI2 B) COS C) CaSO4 D) SF6 E) None of the above contain both ionic and covalent bonds. 5) What is the mass (in kg) of 6.89 × 1025 molecules of CO2? The molar mass of CO2 is 44.01 g/mol. A) 3.85 kg B) 5.04 kg C) 2.60 kg D) 3.03 kg E) 6.39 kg 6) Calculate the mass percent composition of lithium in Li3PO4. A) 26.75 % B) 17.98 % C) 30.72 % D) 55.27 % E) 20.82 % 2) Write the formula for barium nitrite. A) Ba3N2 B) BaNO3 C) BN D) Ba(NO2)2 E) B(NO2)3 7) Determine the molecular formula of a compound that is 49.48% carbon, 5.19% hydrogen, 28.85% nitrogen, and 16.48% oxygen. The molecular weight is 194.19 g/mol. A) C8H12N4O2 B) C4H5N2O C) C8H10N4O2 D) C8H10N2O 3) What is the charge on the Cr ions in Cr2O3? A) 2B) 1+ C) 2+ D) 3+ 4) Calculate the molar mass of Al(C2H3O2)3. A) 86.03 g/mol B) 204.13 g/mol C) 56.00 g/mol D) 258.09 g/mol E) 139.99 g/mol 8) Write a balanced equation to show the reaction of sulfurous acid with lithium hydroxide to form water and lithium sulfite. A) H2SO4 (aq) + LiOH (aq) → H2O (l) + Li2SO4 (aq) B) H2SO3 (aq) + 2 LiOH (aq) → 2 H2O (l) + Li2SO3 (aq) C) HSO3 (aq) + LiOH (aq) → H2O (l) + LiSO3 (aq) D) HSO4 (aq) + LiOH (aq) → H2O (l) + LiSO4 (aq) E) H2S (aq) + 2 LiOH (aq) → 2 H2O (l) + Li2S (aq) Chapter 4 2) How many molecules of H2S are required to form 79.0 g of sulfur according to the following reaction? Assume excess SO2. 1) According to the following balanced reaction, how many moles of HNO3 are formed from 8.44 moles of NO2 if there is plenty of water present? 2 H2S(g) + SO2(g) → 3 S(s) + 2H2O(l) A) 1.48 × 1024 molecules H2S B) 9.89 × 1023 molecules H2S 3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g) C) 5.06 × 1025 molecules H2S D) 3.17 × 1025 molecules H2S A) 2.81 moles HNO3 B) 25.3 moles HNO3 C) 8.44 moles HNO3 D) 5.63 moles HNO3 E) 1.83 moles HNO3 E) 2.44 × 1023 molecules H2S 96 3) Determine the theoretical yield of H2S (in moles) if 4.0 mol Al2S3 and 4.0 mol H2O are reacted according to the following balanced reaction. 9) What volume of 0.305 M AgNO3 is required to react exactly with 155.0 mL of 0.274 M Na2SO4 solution? Hint: You will want to write a balanced reaction. A) 581 mL B) 173 mL C) 345 mL D) 139 mL E) 278 mL Al2S3 (s) + 6 H2O(l) → 2 Al(OH)3 (s) + 3 H2S(g) A) 12 mol H2S B) 4.0 mol H2S C) 18 mol H2S D) 6.0 mol H2S E) 2.0 mol H2S 10) Which of the following solutions will have the highest electrical conductivity? A) 0.045 M Al2(SO4)3 B) 0.050 M (NH4)2CO3 C) 0.10 M LiBr D) 0.10 M NaI E) 0.10 M KF 4) How many moles of LiI are contained in 258.6 mL of 0.0296 M LiI solution? A) 1.31 × 10-3 mol B) 8.74 × 10-3 mol C) 1.14 × 10-3 mol D) 3.67 × 10-3 mol E) 7.65 × 10-3 mol 11) Identify the spectator ions in the following molecular equation. KBr(aq) + AgNO3(aq) → AgBr(s) + KNO3(aq) 5) What is the concentration of magnesium ions in a 0.125 M Mg(NO3)2 solution? A) 0.125 M B) 0.0625 M C) 0.375 M D) 0.250 M E) 0.160 M A) Ag+ and BrB) K+ and NO3C) K+ and BrD) Ag+ and NO3E) There are no spectator ions in this reaction. 6) What volume (in mL) of 0.0887 M MgF2 solution is needed to make 275.0 mL of 0.0224 M MgF2 solution? A) 72.3 mL B) 91.8 mL C) 10.9 mL D) 69.4 mL E) 14.4 mL 12) Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of Na2CO3 and HCl are mixed. A) 2 H+(aq) + CO32-(aq) → H2CO3(s) B) 2 Na+(aq) + CO32-(aq) + 2 H+(aq) + 2 Cl-(aq) → H2CO3(s) + 2 NaCl(aq) C) 2 H+(aq) + CO32-(aq) → H2O(l) + CO2(g) D) 2 Na+(aq) + CO32-(aq) + 2 H+(aq) + 2 Cl-(aq) → H2CO3(s) + 2 Na+(aq) + 2 Cl-(aq) 7) Determine the molarity of a solution formed by dissolving 468 mg of MgI2 in enough water to yield 50.0 mL of solution. A) 0.0297 M B) 0.0337 M C) 0.0936 M D) 0.0107 M E) 0.0651 M E) No reaction occurs. 13) Determine the oxidation state of S in MgSO4. A) -4 B) +2 C) +4 D) +6 E) -2 8) How many ions are present in 30.0 mL of 0.600 M Na2CO3 solution? A) 3.25 × 1022 ions B) 2.17 × 1022 ions C) 1.08 × 1022 ions D) 5.42 × 1022 ions E) 3.61 × 1021 ions 14) What element is undergoing oxidation (if any) in the following reaction? CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) A) O B) H C) C D) both C and H E) None of the elements is undergoing oxidation. 97 Chapter 5 8) Using the graph below, determine the gas that has the lowest density at STP. 1) The atmospheric pressure is 715 mm Hg. What is the pressure in inches of Hg? A) 23.9 in Hg B) 0.940 in Hg C) 48.6 in Hg D) 28.1 in Hg E) 30.0 in Hg 2) A gas occupies 3.33 L at 2.23 atm. What is the volume at 2.50 atm? A) 1.67 L B) 3.73 L C) 2.97 L D) 0.268 L E) 18.6 L 3) A sample of 0.300 moles of nitrogen occupies 0.600 L. Under the same conditions, what number of moles occupies 1.200 L? A) 0.600 moles B) 1.50 moles C) 0.33 moles D) 6.00 moles A) A B) B C) C D) D E) All of the gases have the same density at STP. 9) Define vapor pressure. A) partial pressure of water in a liquid mixture B) partial pressure of water in a gaseous mixture C) condensation of water D) water dissolved in a liquid E) water molecules 4) A syringe initially holds a sample of gas with a volume of 285 mL at 355 K and 1.88 atm. To what temperature must the gas in the syringe be heated/cooled in order to have a volume of 435 mL at 2.50 atm? A) 139 K B) 572 K C) 175 K D) 466 K E) 721 K 10) A mixture of 0.220 moles CO, 0.350 moles H2 and 0.640 moles He has a total pressure of 2.95 atm. What is the pressure of CO? A) 1.86 atm B) 0.649 atm C) 0.536 atm D) 1.54 atm E) 0.955 atm 5) What is the volume of 9.783 x 1023 atoms of He at 9.25 atm and 512K? A) 7.38 L B) 3.69 L C) 1.85 L D) 15.4 L E) 30.8 L 11) The following reaction is used to generate hydrogen gas in the laboratory. If 243 mL of gas is collected at 25 °C and has a total pressure of 745 mm Hg, what mass of hydrogen is produced? A possibly useful table of water vapor pressures is provided below. 6) The density of a gas is 1.43 g/L at 1.00 atm and 273 K. What is the gas? A) Cl2 B) S C) O2 D) Ne Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) T (°C) 20 25 30 7) A compound is found to be 30.45% N and 69.55 % O by mass. If 1.63 g of this compound occupy 389 mL at 0.00°C and 775 mm Hg, what is the molecular formula of the compound? A) NO2 B) N2O C) N4O2 D) N2O5 E) N2O4 P (mm Hg) 17.55 23.78 31.86 A) 0.0196 g H2 B) 0.0717 g H2 C) 0.0190 g H2 D) 0.0144 g H2 E) 0.0449 g H2 98 13) Which of the following statements is TRUE? A) At a given temperature, lighter gas particles travel more slowly than heavier gas particles. B) The smaller a gas particle, the slower it will effuse C) The higher the temperature, the lower the average kinetic energy of the sample. D) At low temperatures, intermolecular forces become important and the pressure of a gas will be lower than predicted by the ideal gas law. E) None of the above statements are true. 12) Determine the volume of O2 (at 1.00 atm and 273 K) formed when 50.0 g of KClO3 decomposes according to the following reaction. 2 KClO3(s) → 2 KCl(s) + 3 O2(g) A) 9.14 L B) 8.22 L C) 12.3 L D) 13.7 L E) 14.6 L Chapter 6 5) According to the following reaction, how much energy is required to decompose 55.0 kg of Fe3O4? 1) Which of the following (with specific heat capacity provided) would show the smallest temperature change upon gaining 200.0 J of heat? A) 50.0 g Al, CAl = 0.903 J/g°C B) 50.0 g Cu, CCu = 0.385 J/g°C C) 25.0 g granite, Cgranite = 0.79 J/g°C D) 25.0 g Au, CAu = 0.128 J/g°C E) 25.0 g Ag, CAg = 0.235 J/g°C Fe3O4(s) → 3 Fe(s) + 2 O2(g) ΔH°rxn = +1118 kJ A) 1.10 × 106 kJ B) 2.38 × 102 kJ C) 2.66 × 105 kJ D) 1.12 × 103 kJ E) 3.44 × 104 kJ 2) Calculate the amount of heat (in kJ) required to raise the temperature of a 79.0 g sample of ethanol from 298.0 K to 385.0 K. The specific heat capacity of ethanol is 2.42 J/g°C. A) 57.0 kJ B) 16.6 kJ C) 73.6 kJ D) 28.4 kJ E) 12.9 kJ 6) Using the following thermochemical equation, determine the amount of heat produced from the combustion of 24.3 g benzene (C6H6). 2 C6H6(l) + 15 O2(g) → 12 CO2(g) + 6 H2O(g) ΔH°rxn = -6278 kJ/mol A) -3910 kJ B) -1950 kJ C) -977 kJ D) -40.1 kJ E) -0.302 kJ 3) Which of the following processes is exothermic? A) the formation of dew in the morning B) the melting of ice C) the chemical reaction in a "cold pack" often used to treat injuries D) the vaporization of water E) None of the above are exothermic. 7) According to the following reaction, how much energy is evolved during the reaction of 32.5 g B2H6 and 72.5 g Cl2? B2H6(g) + 6 Cl2(g) → 2 BCl3(g) + 6 HCl(g) ΔH°rxn = -1396 kJ/mol 4) Using the following equation for the combustion of octane, calculate the heat of reaction for 100.0 g of octane. A) -1640 kJ B) -238 kJ C) -1430 kJ D) -3070 kJ E) -429 kJ 2 C8H18 + 25 O2 → 16 CO2 + 18 H2O ΔH°rxn = -11018 kJ A) -4.82 x 103 kJ/mol B) -4.82 kJ/mol C) -9.64 x 103 kJ/mol D) -1.26 x 104 kJ/mol 99 ΔH 8) According to the following reaction, how much energy is evolved during the reaction of 2.50 L B2H6 and 5.65 L Cl2 (Both gases are initially at STP)? 11) Use the ΔH°f and ΔH°rxn information provided to calculate ΔH°f for IF: IF7(g) + I2(g) → IF5(g) + 2 IF(g) ΔH°rxn = -1396 kJ B2H6(g) + 6 Cl2(g) → 2 BCl3(g) + 6 HCl(g) ΔH°rxn = -1396 kJ/mol A) -58.7 kJ B) -156 kJ C) -215 kJ D) -352 kJ E) -508 kJ IF7(g) IF5(g) 9) Two aqueous solutions are both at room temperature and are then mixed in a coffee cup calorimeter. The reaction causes the temperature of the resulting solution to fall below room temperature. Which of the following statements is TRUE? A) The products have a lower potential energy than the reactants. B) This type of experiment will provide data to calculate ΔErxn. C) The reaction is exothermic. D) Energy is leaving the system during reaction. E) None of the above statements are true. ΔH°f (kJ/mol) -941 -840 A) 101 kJ/mol B) -146 kJ/mol C) -190. kJ/mol D) -95 kJ/mol E) 24 kJ/mol 12) Use the information provided to determine ΔH°rxn for the following reaction: 9) A 100.0 mL sample of 0.300 M NaOH is mixed with a 100.0 mL sample of 0.300 M HNO3 in a coffee cup calorimeter. If both solutions were initially at 35.00°C and the temperature of the resulting solution was recorded as 37.00°C, determine the ΔH°rxn (in units of kJ/mol NaOH) for the neutralization reaction between aqueous NaOH and HCl. Assume 1) that no heat is lost to the calorimeter or the surroundings, and 2) that the density and the heat capacity of the resulting solution are the same as water. A) -55.7 kJ/mol NaOH B) -169 kJ/mol NaOH C) -16.7 kJ/mol NaOH D) -27.9 kJ/mol NaOH E) - 34.4 kJ/mol NaOH CH4(g) + 3 Cl2(g) → CHCl3(l) + 3 HCl(g) ΔH°rxn = ? CH4(g) CHCl3(l) HCl(g) A) -151 kJ B) -335 kJ C) +662 kJ D) +117 kJ E) -217 kJ 10) Use the standard reaction enthalpies given below to determine ΔH°rxn for the following reaction: P4(g) + 10 Cl2(g) → 4PCl5(s) ΔH°rxn = ? Given: PCl5(s) → PCl3(g) + Cl2(g) ΔH°rxn = +157 kJ/mol P4(g) + 6 Cl2(g) → 4 PCl3(g) ΔH°rxn = -1207 kJ/mol ΔH°rxn = -89 kJ A) -1835 kJ B) -1364 kJ C) -1050. kJ D) -1786 kJ E) -2100. kJ 100 ΔH°f (kJ/mol) -75 -134 -92 Chapter 7 1) The vertical height of a wave is called A) wavelength B) amplitude C) frequency D) area E) median 2) Which of the following visible colors of light has the longest wavelength? A) blue B) green C) yellow D) red E) violet 3) Which of the following occur as the energy of a photon increases? A) the frequency decreases. B) the speed increases. C) the wavelength increases D) the wavelength gets shorter. E) None of the above occur as the energy of a photon increases. 4) Calculate the energy of the orange light emitted, per photon, by a neon sign with a frequency of 4.89 × 1014 Hz. A) 3.09 × 10-19 J B) 6.14 × 10-19 J C) 3.24 × 10-19 J D) 1.63 × 10-19 J E) 5.11 × 10-19 J 5) How many photons are contained in a flash of green light (525 nm) that contains 189 kJ of energy? A) 5.67 × 1023 photons B) 2.01 × 1024 photons C) 1.25 × 1031 photons D) 4.99 × 1023 photons E) 7.99 × 1030 photons 6) Calculate the wavelength of a baseball (m = 155 g) moving at 32.5 m/s. A) 7.60 × 10-36 m B) 1.32 × 10-34 m C) 2.15 × 10-32 m D) 2.68 × 10-34 m E) 3.57 × 10-32 m 7) Determine the velocity of a marble (m = 8.66 g) with a wavelength of 3.46 × 10-33 m. A) 45.2 m/s B) 11.3 m/s C) 22.1 m/s D) 38.8 m/s E) 52.9 m/s 8) Choose the transition (in a hydrogen atom) below that represents the absorption of the shortest wavelength photon. A) n = 1 to n = 2 B) n = 2 to n = 3 C) n = 4 to n = 5 D) n = 6 to n = 3 E) n = 3 to n = 1 9) Which of the following statements is TRUE? A) We can sometimes know the exact location and speed of an electron at the same time. B) All orbitals in a given atom are roughly the same size. C) Since electrons have mass, we must always consider them to have particle properties and never wavelike properties. D) Atoms are roughly spherical because when all of the different shaped orbitals are overlapped, they take on a spherical shape. E) All of the above are true. 10) Identify the correct values for a 2p sublevel. A) n = 3, l = 1, ml = 0 B) n = 2, l = 1, ml = -2 C) n = 1, l = 0, ml = 0 D) n = 2, l = 1, ml = 0 E) n = 4, l = -1, ml = -2 11) Determine the energy change associated with the transition from n=2 to n=5 in the hydrogen atom. A) -2.18 × 10-19 J B) +6.54 × 10-19 J C) +4.58 × 10-19 J D) -1.53 × 10-19 J E) +3.76 × 10-19 J 12) Determine the end (final) value of n in a hydrogen atom transition, if the electron starts in n = 4 and the atom emits a photon of light with a wavelength of 486 nm. A) 1 B) 5 C) 3 D) 4 E) 2 13) What is the maximum number of f orbitals that are possible? A) 1 B) 3 C) 7 D) 5 E) 9 101 Chapter 8 6) Place the following elements in order of decreasing atomic radius. 1) Choose the orbital diagram that represents the ground state of N. A) Xe D) Ar A) Ar > Xe > Rb B) Xe > Rb > Ar C) Ar > Rb > Xe D) Rb > Xe > Ar E) Rb > Ar > Xe B) C) Rb 7) Of the following, which atom has the smallest atomic radius? A) K B) As C) Rb D) Sb 8) Place the following in order of increasing radius. Ca2+ E) 2) The element that corresponds to the electron configuration 1s22s22p6 is __________. A) sodium B) magnesium C) lithium D) beryllium E) neon 3) The complete electron configuration of argon, element 18, is __________. A) 1s22s22p63s23p6 B) 1s22s22p103s23p2 C) 1s42s42p63s4 D) 1s42s42p10 E) 1s62s 62p23s4 4) Give the ground state electron configuration for Se. A) [Ar]4s23d104p4 B) [Ar]4s24d104p4 C) [Ar]4s23d104p6 D) [Ar]4s23d10 E) [Ar]3d104p4 S2- Cl- A) Ca2+ < Cl- < S2B) Cl- < Ca2+ < S2C) S2- < Cl- < Ca2+ D) Ca2+ < S2- < ClE) Cl- < S2- < Ca2+ 9) Which reaction below represents the second ionization of Sr? A) Sr(g) → Sr+ (g) + eB) Sr2+ (g) + e- → Sr- (g) C) Sr+ (g) + e- → Sr(g) D) Sr- (g) + e- → Sr2- (g) E) Sr+ (g) → Sr2+ (g) + e10) Choose the ground state electron configuration for Ti2+. A) [Ar]3d2 B) [Ar]4s2 C) [Ar]4s23d2 D) [Ar]4s23d4 E) [Ar]3d4 5) How many unpaired electrons are present in the ground state Kr atom? A) 1 B) 2 C) 0 D) 3 E) 5 102 Chapter 9 7) Use the data given below to construct a Born-Haber cycle to determine the lattice energy of CaO. 1) Which of the following represent the Lewis structure for N? A) Ca(s) → Ca(g) Ca(g) → Ca+(g) + eCa+ (g) → Ca2+ (g) + e2 O(g) → O2(g) O(g) + e- → O- (g) O- (g) + e- → O2- (g) B) C) D) E) 2) Which of the following statements is TRUE? A) An ionic bond is much stronger than most covalent bonds. B) An ionic bond is formed through the sharing of electrons. C) Ionic compounds at room temperature typically conduct electricity. D) Once dissolved in water, ionic compounds rarely conduct electricity. E) None of the above are true. 3) Identify the compound with the highest magnitude of lattice energy. A) NaCl B) KCl C) LiCl D) CsCl Ca(s) + A) -3414 kJ B) +1397 kJ C) -2667 kJ D) +3028 kJ E) -2144 kJ 8) Choose the bond below that is most polar. A) H-I B) H-Br C) H-F D) H-Cl E) C-H 9) Choose the best Lewis structure for CH2Cl2. A) 4) Choose the compound below that should have the highest melting point according to the ionic bonding model. A) AlN B) MgO C) NaF D) CaS E) RbI B) C) 5) Give the complete electronic configuration for Ca2+. A) 1s22s22p63s24p6 B) 1s22s22p63s23p6 C) 1s22s22p63s23p5 D) 1s22s23p64s25p6 E) 1s22s2p63s2p6 D) E) 6) Place the following elements in order of decreasing electronegativity. S Cl 1 O2(g) → CaO(s) 2 ΔH°(kJ/mol) 193 590 1010 -498 -141 878 -635 Se A) Se > S > Cl B) Cl > Se > S C) Se > Cl > S D) S > Cl > Se E) Cl > S > Se 103 10) Choose the best Lewis structure for SF4. A) 14) Choose the best Lewis structure for SO42 . A) B) B) C) C) D) D) E) E) 11) Which of the following resonance structures for OCN- will contribute most to the correct structure of OCN-? A) O(2 lone pairs)=C=N (2 lone pairs) B) O(1 lone pair)≡C-N(3 lone pairs) C) O(1 lone pair)=C(2 lone pairs)=N(1 lone pair) D) O(3 lone pairs)-C≡N(with 1 lone pair) E) They all contribute equally to the correct structure of OCN-. 12) Which of the following elements can form compounds with an expanded octet? A) N B) Br C) F D) Be E) None of the above can form compounds with an expanded octet. 13) Place the following in order of decreasing bond length. H-F H-I H-Br A) H-F > H-Br > H-I B) H-I > H-F > H-Br C) H-I > H-Br > H-F D) H-Br > H-F > H-I E) H-F > H-I > H-Br 15) Use the bond energies provided to estimate ΔH°rxn for the reaction below. CH3OH(l) + 2 O2(g) → CO2(g) + 2 H2O(g) Bond Bond Energy (kJ/mol) C-H 414 C-O 360 C=O 799 O=O 498 O-H 464 A) +473 kJ B) -91 kJ C) -486 kJ D) -392 kJ E) +206 kJ 16) Which compound has the highest carbon-carbon bond strength? A) CH3CH3 B) CH2CH2 C) HCCH D) all bond strengths are the same 104 ΔH Chapter 10 1) Give the approximate bond angle for a molecule with a trigonal planar shape. A) 109.5° B) 180° C) 120° D) 105° E) 90° 2) Give the approximate bond angle for a molecule with an octahedral shape. A) 109.5° B) 180° C) 120° D) 105° E) 90° 3) Determine the electron geometry (eg) of CO32-. A) eg=linear B) eg=tetrahedral C) eg=trigonal planar D) eg=trigonal bipyramidal E) eg=octahedral 4) Determine the electron geometry (eg) of NCl3. A) eg=linear B) eg=tetrahedral C) eg=trigonal planar D) eg=trigonal bipyramidal E) eg=octahedral 5) Determine the electron geometry (eg) of XeF4. A) eg=linear B) eg=tetrahedral C) eg=trigonal planar D) eg=trigonal bipyramidal E) eg=octahedral 6) Place the following in order of increasing F-A-F bond angle, where A represents the central atom in each molecule. PF3 OF2 PF43- A) PF3 < OF2 < PF43B) OF2 < PF3 < PF43C) OF2 < PF43- < PF3 D) PF43- < OF2 < PF3 E) PF43- < PF3 < OF2 7) A molecule containing a central atom with sp3d hybridization has a(n) __________ electron geometry. A) tetrahedral B) linear C) octahedral D) trigonal planar E) trigonal bipyramidal 8) Place the following in order of decreasing X-A-X bond angle, where A represents the central atom and X represents the outer atoms in each molecule. N 2O NCl3 NO2- A) NCl3 > NO2- > N2O B) NO2- > N2O > NCl3 C) N2O > NO2- > NCl3 D) NCl3 > N2O > NO2E) N2O > NCl3 > NO29) How many of the following molecules are polar? XeCl2 COF2 PCl4F SF6 A) 0 B) 3 C) 1 D) 2 E) 4 10) Determine the electron geometry and polarity of SF6 . A) eg=trigonal bipyramidal, nonpolar B) eg=tetrahedral, polar C) eg=trigonal bipyramidal, polar D) eg=octahedral, nonpolar E) eg=octahedral, polar 11) Determine the electron geometry and polarity of HBrO2 . A) eg=trigonal bipyramidal, nonpolar B) eg=octahedral, nonpolar C) eg=tetrahedral, polar D) eg=tetrahedral, nonpolar E) eg=linear, polar 12) Choose the compound below that contains at least one polar covalent bond, but is nonpolar. A) GeH2Br2 B) SCl2 C) AsCl5 D) CF2Cl2 E) All of the above are nonpolar and contain a polar covalent bond. 13) Describe a pi bond. A) side by side overlap of p orbitals B) end to end overlap of p orbitals C) s orbital overlapping with the end of a p orbital D) overlap of two s orbitals E) p orbital overlapping with a d orbital 105 14) Use the molecular orbital diagram shown to determine which of the following are paramagnetic. 17) Use the molecular orbital diagram shown to determine which of the following is most stable. A) B22+ B) B22- A) C22B) N22- C) N22+ D) C22- C) B2 D) C22E) B22- E) B2 15) Determine the electron geometry (eg) of the underlined carbon in CH3CN. A) eg=tetrahedral B) eg=linear C) eg=trigonal planar D) eg=trigonal bipyramidal E) eg=octahedral 16) Draw the Lewis structure for H3O+. What is the hybridization on the O atom? A) sp B) sp3 C) sp2 D) sp3d E) sp3d2 18) Draw the molecular orbital diagram shown to determine which of the following is paramagnetic. A) B22+ B) B22C) N22D) C22E) B2 19) Draw the Lewis structure for BrF5. What is the hybridization on the Br atom? A) sp3d2 B) sp3d C) sp3 D) sp2 E) sp 106 20) How many of the following molecules have sp3 hybridization on the central atom? XeCl4 CH4 SF4 21) How many of the following molecules have sp3d2 hybridization on the central atom? C 2H 2 SeCl6 A) 0 B) 4 C) 3 D) 2 E) 1 Chapter 11 3) What type of intermolecular force causes the dissolution of NaCl in water? A) hydrogen bonding B) dipole-dipole forces C) ion-dipole force D) dispersion forces E) none of the above 4) Choose the molecule or compound that exhibits dipole-dipole forces as its strongest intermolecular force. A) H2 B) SO2 C) NH3 D) CF4 E) BCl3 IF5 AsCl5 A) 1 B) 3 C) 0 D) 2 E) 4 6) Place the following compounds in order of increasing strength of intermolecular forces. 1) Which of the following statements is TRUE? A) Intermolecular forces are generally stronger than bonding forces. B) The potential energy of molecules decrease as they get closer to one another. C) Energy is given off when the attraction between two molecules is broken. D) Increasing the pressure on a solid usually causes it to become a liquid. E) None of the above are true. 2) What is the strongest type of intermolecular force present in CHF3? A) ion-dipole B) dispersion C) hydrogen bonding D) dipole-dipole E) none of the above XeF4 CH4 CH3CH2CH3 CH3CH3 A) CH3CH2CH3 < CH4 < CH3CH3 B) CH3CH2CH3 < CH3CH3 < CH4 C) CH3CH3 < CH4 < CH3CH2CH3 D) CH4 < CH3CH2CH3 < CH3CH3 E) CH4 < CH3CH3 < CH3CH2CH3 7) Choose the pair of substances that are most likely to form a homogeneous solution. A) C6H14 and C10H20 B) LiBr and C5H12 C) N2O4 and NH4Cl D) C6H14 and H2O E) None of the pairs above will form a homogeneous solution. 8) Which of the following statements is FALSE? A) The rate of vaporization increases with increasing surface area. B) The rate of vaporization increases with decreasing strength of intermolecular forces. C) The rate of vaporization increases with increasing temperature. D) Molecules with hydrogen bonding are more volatile than compounds with dipole-dipole forces. E) None of the above are false. 9) Place the following substances in order of increasing boiling point. CH3CH2OH Ar CH3OCH3 A) Ar < CH3OCH3 < CH3CH2OH B) CH3CH2OH < Ar < CH3OCH3 C) CH3CH2OH < CH3OCH3 < Ar D) CH3OCH3 < Ar < CH3CH2OH E) Ar < CH3CH2OH < CH3OCH3 5) Identify the compound that does not have hydrogen bonding. A) (CH3)3N B) H2O C) CH3OH D) HF E) CH3NH2 107 10) Which of the following substances would you predict to have the highest ΔHvap? A) Xe B) C6H6 C) SiF4 D) Br2 E) N2 11) How much energy is required to vaporize 158 g of butane (C4H10) at its boiling point, if its ΔHvap is 24.3 kJ/mol? A) 15.1 kJ B) 66.1 kJ C) 89.4 kJ D) 11.2 kJ E) 38.4 kJ 12) Determine the normal boiling point of a substance whose vapor pressure is 55.1 mm Hg at 35°C and has a ΔHvap of 32 .1 kJ/mol. A) 255 K B) 368 K C) 412 K D) 390. K E) 466 K 13) Ethanol (C2H5OH) melts at -114°C. The enthalpy of fusion is 5.02 kJ/mol. The specific heats of solid and liquid ethanol are 0.97 J/gK and 2.3 J/gK, respectively. How much heat (kJ) is needed to convert 25.0 g of solid ethanol at -135°C to liquid ethanol at -50°C? A) 207.3 kJ B) -12.7 kJ C) 6.91 kJ D) 4192 kJ E) 9.21 kJ Chapter 12 1) Dynamic equilibrium can be defined as A) rate of dissolution = rate of deposition B) rate of dissolution < rate of deposition C) rate of dissolution > rate of deposition D) rate of bubbling > rate of dissolving E) rate of evaporating > rate of condensing 2) What mass (in g) of NH3 must be dissolved in 475 g of methanol to make a 0.250 m solution? A) 2.02 g B) 4.94 g C) 1.19 g D) 8.42 g E) 1.90 g 14) Based on the figure above, the boiling point of water under an external pressure of 0.316 atm is __________°C. A) 70 B) 40 C) 60 D) 80 E) 90 3) Determine the molality of a solution prepared by dissolving 0.500 moles of CaF2 in 11.5 moles H2O. A) 1.88 m B) 4.35 m C) 5.31 m D) 4.14 m E) 2.41 m 4) A solution is prepared by dissolving 98.6 g of NaCl in enough water to form 875 mL of solution. Calculate the mass % of the solution if the density of the solution is 1.06 g/mL. A) 11.3% B) 12.7% C) 9.4% D) 10.6% E) 11.9% 108 5) Determine the freezing point of a solution that contains 78.8 g of naphthalene (C10H8, molar mass = 128.16 g/mol) dissolved in 722 mL of benzene (d = 0.877 g/mL). Pure benzene has a melting point of 5.50°C and a freezing point depression constant of -4.90°C/m. A) 4.76°C B) 4.17°C C) 0.74°C D) 1.33°C E) 1.68°C 8) A 150.0 mL sample of an aqueous solution at 25°C contains 15.2 mg of an unknown nonelectrolyte compound. If the solution has an osmotic pressure of 8.44 torr, what is the molar mass of the unknown compound? A) 223 g/mol B) 294 g/mol C) 341 g/mol D) 448 g/mol E) 195 g/mol 9) Choose the aqueous solution that has the highest boiling point. These are all solutions of nonvolatile solutes and you should assume ideal van't Hoff factors where applicable. A) 0.100 m NaNO3 B) 0.100 m Li2SO4 C) 0.200 m C3H8O3 D) 0.060 m Na3PO4 E) They all have the same boiling point. 6) Calculate the boiling point of a solution of 500.0 g of ethylene glycol (C2H6O2) dissolved in 500.0 g of water. Kf = 1.86°C/m and Kb = 0.512°C/m. Use 100°C as the boiling point of water. A) 108°C B) 92°C C) 130°C D) 70°C E) 8.3°C 7) Determine the boiling point of a solution that contains 78.8 g of naphthalene (C10H8, molar mass = 128.16 g/mol) dissolved in 722 mL of benzene (d = 0.877 g/mL). Pure benzene has a boiling point of 80.1°C and a boiling point elevation constant of 2.53°C/m. A) 2.2°C B) 2.5°C C) 82.3°C D) 80.4°C E) 82.6°C 10) Choose the aqueous solution below with the lowest freezing point. These are all solutions of nonvolatile solutes and you should assume ideal van't Hoff factors where applicable. A) 0.075 m NaI B) 0.075 m (NH4)3PO4 C) 0.075 m NaBrO4 D) 0.075 m LiCN E) 0.075 m KNO2 109