Thursday, May 23, 2002

Honors Chemistry Semester 2 Final Review

1. Which of the following frequencies corresponds to light with the longest wavelength?
A) 3.00 ´ 1013 s–1
B) 9.12 ´ 1012 s–1
C) 3.20 ´ 109 s–1
D) 8.50 ´ 1020 s–1
E) 4.12 ´ 105 s–1

2. Calculate the lattice energy for LiF(s) given the following:
sublimation energy for Li(s) +166 kJ/mol
DHf for F(g) +77 kJ/mol
first ionization energy of Li(g) +520. kJ/mol
electron affinity of F(g) –328 kJ/mol
enthalpy of formation of LiF(s) –617 kJ/mol
A) 285 kJ/mol
B) –1047 kJ/mol
C) –650. kJ/mol
D) 800. kJ/mol
E) None of these

3. Name the following:

A) 3,3-dimethylpentane
B) 2,2-diethylpropane
C) n-heptane
D) 2-methyl-2-ethylbutane

4. Which of the following statements is true?
A) The exact location of an electron can be determined if we know its energy.
B) In the buildup of atoms, electrons occupy the 4f orbitals before the 6s orbitals.
C) Only three quantum numbers are needed to uniquely describe an electron.
D) Ni has 2 unpaired electrons in its 3d orbitals.
E) An electron in a 2s orbital can have the same n, l, and ml quantum numbers as an electron in a 3s orbital.

5. A solution of hydrogen peroxide is 30.0% H2O2 by mass and has a density of 1.11 g/cm3. The molarity of the solution is:
A) 8.82 M
B) 0.980 M
C) 9.79 M
D) 7.94 M
E) none of these

6. Consider the reaction X ------à Y + Z
Which of the following is a possible rate law?
A) Rate = k[X]
B) Rate = k[Z]
C) Rate = k[X][Y]
D) Rate = k[Y][Z]
E) Rate = k[Y]

7. On a relative basis, the weaker the intermolecular forces in a substance,
A) the higher its melting point.
B) the greater its vapor pressure at a particular temperature.
C) the more it deviates from ideal gas behavior.
D) the greater its heat of vaporization.
E) none of these

8. The vapor pressure of water at 25.0°C is 23.8 torr. Determine the mass of glucose (molar mass = 180 g/mol) needed to add to 500.0 g of water to change the vapor pressure to 23.1 torr.
A) 72 g
B) 103 g
C) 115 g
D) 36 g
E) 152 g

9. One of the ingredients on a margarine container is listed as "polyunsaturated corn oil." This means that
A) all the carbon-carbon bonds are triple bonds.
B) many of the carbon-carbon bonds are multiple bonds.
C) all the carbon bonds in the oil are single bonds.
D) many of the polymer bonds are unsaturated.
E) none of these

10. Assume that the enthalpy of fusion of ice is 6020 J/mol and does not vary appreciably over the temperature range 270-290 K. If one mole of ice at 0°C is melted by heat supplied from surroundings at 280 K, what is the entropy change in the surroundings, in J/K?
A) 0.0
B) –21.5
C) +21.5
D) –22.0
E) +22.0

11. Which of the following has the largest radius?
A) Ar
B) K+
C) Cl–
D) Ca2+
E) S2–

12. In the molecule C2H4 the valence orbitals of the carbon atoms are assumed to be
A) sp hybridized.
B) sp2 hybridized.
C) not hybridized.
D) dsp hybridized.
E) sp3 hybridized.

13. Choose the statement that best describes the PbCl4 molecule in the gas phase.
A) The bond angles are all about 109°.
B) The molecule is polar.
C) The molecule has a dipole moment.
D) The bonds are nonpolar.
E) a, b, and c

14. Alpha particles beamed at thin metal foil may
A) be reflected by direct contact with nuclei.
B) pass directly through without changing direction and be reflected by direct contact with nuclei.
C) be slightly diverted by attraction to electrons.
D) pass directly through without changing direction, be slightly diverted by attraction to electrons, and be reflected by direct contact with nuclei.
E) pass directly through without changing direction.

15. Which of the following processes increases the atomic number by 1?
A) beta-particle production
B) positron production
C) alpha-particle production
D) electron capture
E) gamma-ray production

16. How many electrons in an atom can have the quantum numbers n = 3, l = 2?
A) 18
B) 5
C) 6
D) 2
E) 10

17. Determine the molarity of a solution of the weak acid HClO2 (Ka = 1.10 ´ 10–2) if it has a pH of 1.25.
A) 1.23 M
B) 1.52 M
C) 3.17 M
D) 0.819 M
E) 0.287 M

18. How many f orbitals have the value n = 3?
A) 5
B) 7
C) 1
D) 3
E) 0

19. Concentrated nitric acid is a solution that is 70% HNO3, by mass. The density of this acid is 1.42 g/cm3. What is the molarity of this acid?
A) 0.16 M
B) 15.8 M
C) 6.8 M
D) 12.4 M
E) 0.06 M

20. Which of the following is incorrect?
A) The continuous spectrum of hydrogen contains only four discrete colors.
B) All matter displays both particle and wavelike characteristics.
C) The lowest possible energy state of a molecule or atom is called its ground state.
D) Diffraction produces both constructive and destructive interference.
E) Niels Bohr developed a quantum model for the hydrogen atom.

21. Fresh rainwater or surface water contains enough tritium to show 5.5 decompositions per minute per 100. g of water. Tritium has a half-life of 12.3 years. You are asked to check a vintage wine claimed to have been produced in 1946. How many decompositions per minute should you expect to observe in 100. g of that wine?
A) 1.7
B) 0.35
C) 49
D) 0.035
E) 0.49

22. Consider the reaction:

at constant temperature. Initially a container is filled with pure SO3(g) at a pressure of 2 atm, after which equilibrium is allowed to be reached. If y is the partial pressure of O2 at equilibrium, the value of Kp is:
E) none of these

23. At 40°C, heptane has a vapor pressure of 92.0 torr and octane has a vapor pressure of 31.2 torr. Assuming ideal behavior, what is the vapor pressure of a solution that contains twice as many moles of heptane as octane?
A) 71.7 torr
B) 61.6 torr
C) 51.5 torr
D) 76.8 torr
E) none of these

Use the following to answer question 24:

Solutions of benzene and toluene obey Raoult's law. The vapor pressures at 20°C are:
benzene, 76 torr; toluene, 21 torr.

24. The vapor pressure of water at 90°C is 0.692 atm. What is the vapor pressure (in atm) of a solution made by dissolving 1.00 mole of CsF(s) in 1.00 kg of water? Assume that Raoult's law applies.
A) 0.668 atm
B) 0.680 atm
C) 0.692 atm
D) 0.656 atm
E) none of these

25. In the following nuclear equation, identify the missing product:


26. Green light has a wavelength of 5.50 ´ 102 nm. The energy of a photon of green light is
A) 5.45 ´ 1012 J
B) 2.17 ´ 105 J
C) 3.64 ´ 10–38 J
D) 3.61 ´ 10–19 J
E) 1.09 ´ 10–27 J

27. Consider the reaction 2H2 + O2 ------à 2H2O
What is the ratio of the initial rate of the appearance of water to the initial rate of disappearance of oxygen?
A) 1 : 2
B) 3 : 2
C) 1 : 1
D) 2 : 1
E) 2 : 2

Use the following to answer questions 28-33:

A general reaction written as 2A + 2B ------à C + 2D is studied and yields the following data:
[A]0 [B]0 Initial D[C]/Dt
0.100 M 0.100 M 0.000040 mol/L × s
0.200 M 0.100 M 0.000160 mol/L × s
0.100 M 0.200 M 0.000040 mol/L × s

28. What are the proper units for the rate constant for the reaction?
A) L3 mol–3 s–1
B) L2 mol–2 s–1
C) L mol–1 s–1
D) s–1
E) mol L–1 s–1

29. What is the order of the reaction with respect to A?
A) 2
B) 0
C) 1
D) 3
E) 4

30. What is the overall order of the reaction?
A) 0
B) 3
C) 1
D) 4
E) 2

31. What is the order of the reaction with respect to B?
A) 0
B) 2
C) 1
D) 4
E) 3

32. For the first of the reactions in the table of data, how many seconds would it take for [A] to decrease to 0.050 M?
A) 2500
B) 1200
C) 250
D) 1700
E) 170

33. What is the numerical value of the rate constant?
A) 0.0160
B) 0.000160
C) 0.0040
D) 4.0 ´ 10–7
E) 0.000040

34. What is the wavelength of light that is emitted when an excited electron in the hydrogen atom falls from n=5 to n=2?
A) 4.34 × 10–7 m
B) 6.50 × 10–7 m
C) 5.12 × 10–7 m
D) 5.82 × 10–7 m
E) none of these

35. Which of the following has the smallest radius?
A) K+
B) Cl–
C) Al3+
D) Ar
E) Cr

36. Which of the following concentration measures will change in value as the temperature of a solution changes?
A) molarity
B) molality
C) mass percent
D) mole fraction
E) all of these

37. The term "proof" is defined as twice the percent by volume of pure ethanol in solution. Thus, a solution that is 95% (by volume) ethanol is 190 proof. What is the molarity of ethanol in a 92 proof ethanol/water solution?
density of ethanol = 0.80 g/cm3
density of water = 1.0 g/cm3
mol. wt. of ethanol = 46
A) 0.92 M
B) 17 M
C) 0.80 M
D) 0.46 M
E) 8.0 M

Use the following to answer questions 38-39:

Consider the chemical system
K = 4.6 ´ 109 L/mol.

38. How do the equilibrium concentrations of the reactants compare to the equilibrium concentration of the product?
A) They are much smaller.
B) They are much bigger.
C) You can't tell from the information given.
D) They have to be exactly equal.
E) They are about the same.

39. If the concentration of the product were to double, what would happen to the equilibrium constant?
A) It would quadruple its value.
B) It would depend on the initial conditions of the product.
C) It would double its value.
D) It would not change its value.
E) It would become half its current value.

40. Which of the following contains a pi bond?
A) C2H6
D) CCl4
E) PF2

41. The rate constant for the beta decay of thorium-234 is 2.88 ´ 10–2/day. What is the half-life of this nuclide?
A) 1.22 days
B) 0.693 days
C) 101 days
D) 53.1 days
E) 24.1 days

42. 15.0 mL of 0.50 M HCl is added to a 100.-mL sample of 0.200 M HNO2
(Ka for HNO2 = 4.0 ´ 10–14). What is the equilibrium concentration of NO2– ions?
A) 7.5 ´ 10–14 M
B) 1.1 ´ 10–13 M
C) 6.5 ´ 10–12 M
D) 1.7 ´ 10–11 M
E) none of these

43. The value of the equilibrium constant, K, is dependent on
I. The temperature of the system.
II. The nature of the reactants and products.
III. The concentration of the reactants.
IV. The concentration of the products.
A) I, II
D) It is dependent on three of the above choices.
E) It is not dependent on any of the above choices.

44. The conjugate base of a weak acid is
A) a weak base
B) a strong acid
C) a strong base
D) a weak acid
E) none of these

45. A 10.0-g sample of solid NH4Cl is heated in a 5.00-L container to 900°C. At equilibrium the pressure of NH3(g) is 1.20 atm.

The equilibrium constant, Kp, for the reaction is:
A) 1.20
B) 2.40
C) 31.0
D) 1.44
E) none of these

46. What element is the major component of bone?
A) phosphorus
B) potassium
C) calcium
D) carbon
E) oxygen

47. The molecules in a sample of solid SO2 are attracted to each other by a combination of
A) London forces and H-bonding.
B) covalent bonding and dipole-dipole interactions.
C) H-bonding and ionic bonding.
D) London forces and dipole-dipole interactions.
E) none of these

48. Consider the following Lewis structure

Which statement about the molecule is false?
A) There are some H–C–H bond angles of about 109° in the molecule.
B) This molecule contains 28 valence electrons.
C) There are 10 sigma and 2 pi bonds.
D) Oxygen is sp3 hybridized.
E) C–2 is sp2 hybridized with bond angles of 120°.

49. When 0.800 g of NH4NO3 was added to 150.0 g of water in a styrofoam cup, the temperature dropped by 0.413°C. The heat capacity of H2O is 4.18 J/g°C. Assume the specific heat of the solution equals that of pure H2O and that the calorimeter neither absorbs nor leaks heat. The molar heat of solution of solid NH4NO3 is:
A) +2.60 kJ/mol
B) +260. J/mol
C) –2.60 kJ/mol
D) 26.0 kJ/mol
E) none of these

50. Nuclides with too many neutrons to be in the band of stability are most likely to decay by what mode?
A) b–
B) electron capture
C) b+
D) a
E) fission

51. Name the following:

A) 3-methylpentane
B) 1,2,3-trimethylpropane
C) methyl-diethylmethane
D) n-hexane
E) isohexane

52. Which of the following is not true for a solution at 25°C that has a hydroxide concentration of 2.5 ´ 10–6 M?
A) The [H] is 4 ´ 10–9 M.
B) The solution is acidic.
C) The solution is basic.
D) The Kw is independent of what the solution contains.
E) Kw = 1 ´ 10–14

53. Tabulated below are initial rate data for the reaction
2Fe(CN)63– + 2I– ------à 2Fe(CN)64– + I2
Run [Fe(CN)63–]0 [I–]0 [Fe(CN)64-]0 [I2]0 Rate (M/s)
1 0.01 0.01 0.01 0.01 1 ´ 10–5
2 0.01 0.02 0.01 0.01 2 ´ 10–5
3 0.02 0.02 0.01 0.01 8 ´ 10–5
4 0.02 0.02 0.02 0.01 8 ´ 10–5
5 0.02 0.02 0.02 0.02 8 ´ 10–5
The experimental rate law is:
A) k[Fe(CN)63–][I–] [Fe(CN)64–]
B) k[Fe(CN)63–)]2[I–]
C) k[Fe(CN)63–]2[I–][Fe(CN)64–][I2]
D) k[Fe(CN)63–]2[I–]2[Fe(CN)64–]2[I2]
E) k[Fe(CN)63–][I–]2

54. Electron capture transforms into what nuclide?

55. A polypeptide is
A) an addition polymer of amino acids.
B) a polymer of sugar molecules.
C) a part of nucleic acids.
D) a condensation polymer of amino acids.
E) none of these

56. An unstable isotope of rhenium, 191Re, has a half-life of 9.8 minutes and is a beta producer. What is the other product of the reaction?
A) 190W
B) 190Os
C) 192Pt
D) 191Os
E) 191W

57. Consider the reaction whose K = 54.8 at 425°C. If an equimolar mixture of reactants gives the concentration of the product to be 0.50 M at equilibrium, determine the concentration of the hydrogen.
A) 1.6 ´ 10–4 M
B) 4.6 ´ 10–3 M
C) 9.6 ´ 10–2 M
D) 6.8 ´ 10–2 M
E) 1.2 ´ 10–3 M

58. Using the following data reactions
DH° (kJ)
H2(g) + Cl2(g) ------à 2HCl(g) –184
H2(g) ------à 2H(g) 432
Cl2(g) ------à 2Cl(g) 239
calculate the energy of an H-Cl bond.
A) 518 kJ
B) 856 kJ
C) 428 kJ
D) 326 kJ
E) 770 kJ

59. A mixture of hydrogen and chlorine remains unreacted until it is exposed to ultraviolet light from a burning magnesium strip. Then the following reaction occurs very rapidly:
H2(g) + Cl2(g) ------à 2HCl(g) DG = –45.54 kJ
DH = –44.12 kJ
DS = -4.76 J/K
Select the statement below that best explains this behavior.
A) The reaction has a small equilibrium constant.
B) The reaction is spontaneous, but the reactants are kinetically stable.
C) The ultraviolet light raises the temperature of the system and makes the reaction more favorable.
D) The reactants are thermodynamically more stable than the products.
E) The negative value for DS slows down the reaction.

60. For the equilibrium that exists in an aqueous solution of nitrous acid (HNO2, a weak acid), the equilibrium constant expression is:
E) none of these

61. Tabulated below are initial rate data for the reaction
Run [Fe(CN)63–]0 [I–]0 [Fe(CN)64–]0 [I2]0 Rate (M/s)
1 0.01 0.01 0.01 0.01 1 ´ 10–5
2 0.01 0.02 0.01 0.01 2 ´ 10–5
3 0.02 0.02 0.01 0.01 8 ´ 10–5
4 0.02 0.02 0.02 0.01 8 ´ 10–5
5 0.02 0.02 0.02 0.02 8 ´ 10–5
The value of k is:
A) 103 M–3 s–1
B) 107 M–5 s–1
C) 50 M–2 s–1
D) 10 M–2 s–1
E) none of these

62. For the reaction below, Kp = 1.16 at 800°C.

If a 20.0-gram sample of CaCO3 is put into a 10.0-liter container and heated to 800°C, what percent of the CaCO3 will react to reach equilibrium?
A) 65.9%
B) 14.6%
C) 100.0%
D) 34.1%
E) none of these

63. Which statement about hydrogen bonding is true?
A) Hydrogen bonding arises from the dipole moment created by the unequal sharing of electrons within certain covalent bonds within a molecule.
B) Hydrogen bonding of solvent molecules with a solute will not affect the solubility of the solute.
C) The hydrogen bonding capabilities of water molecules cause CH3CH2CH2CH3 to be more soluble in water than CH3OH.
D) Hydrogen bonding interactions between molecules are stronger than the covalent bonds within the molecule.
E) Hydrogen bonding is the intermolecular attractive forces between two hydrogen atoms in solution.

64. The pH of a solution at 25°C in which [OH–] = 3.4 ´ 10–5 M is:
A) 9.5
B) 4.5
C) 6.3
D) 10.5
E) none of these

65. The reaction below occurs in basic solution. In the balanced equation, what is the sum of the coefficients?
Zn + NO3– ------à Zn(OH)42– + NH3
A) 23
B) 12
C) 27
D) 19
E) 15

66. A solution contains 0.250 M HA (Ka = 1.0 × 10–6) and 0.45 M NaA. What is the pH after 0.10 mole of HCl is added to 1.00 L of this solution?
A) 6.00
B) 3.23
C) 10.77
D) 10.83
E) 3.17

67. A solution containing 296.6 g of Mg(NO3)2 per liter has a density of 1.114 g/mL. The molarity of the solution is:
A) 2.000 M
B) 6.001 M
C) 1.805 M
D) 2.446 M
E) none of these

68. Which energy conversion shown below takes place in a galvanic cell?
A) mechanical to electrical
B) mechanical to chemical
C) chemical to electrical
D) chemical to mechanical
E) electrical to chemical

69. The reaction

has Kp = 45.9 at 763 K. A particular equilibrium mixture at that temperature contains gaseous HI at a partial pressure of 4.00 atm and hydrogen gas at a partial pressure of 0.200 atm. What is the partial pressure of I2?
A) 14.3 atm
B) 0.436 atm
C) 1.74 atm
D) 0.200 atm
E) 0.574 atm

70. In an investigation of the electronic absorption spectrum of a particular element, it is found that a photon having l = 500 nm provides just enough energy to promote an electron from the second quantum level to the third. From this information, we can deduce
A) the energy of the n = 3 level.
B) the difference in energies between n = 2 and n = 3.
C) the sum of the energies of n = 2 and n = 3.
D) the energy of the n = 2 level.
E) all of these

71. You have solutions of 0.200 M HNO2 and 0.200 M KNO2 (Ka for HNO2 = 4.00 ´ 10–4). A buffer of pH 3.000 is needed. What volumes of HNO2 and KNO2 are required to make 1 liter of buffered solution?
A) 286 mL HNO2; 714 mL KNO2
B) 587 mL HNO2; 413 mL KNO2
C) 413 mL HNO2; 587 mL KNO2
D) 500 mL of each
E) 714 mL HNO2; 286 mL KNO2

72. What combination of substances will give a buffered solution that has a pH of 5.05? (Assume each pair of substances is dissolved in 5.0 L of water.)
(Kb for NH3 = 1.8 ´ 10–5; Kb for C5H5N = 1.7 ´ 10–9)
A) 1.0 mole NH3 and 1.5 mole NH4Cl
B) 1.5 mole C5H5N and 1.0 mole C5H5NHCl
C) 1.5 mole NH3 and 1.0 mole NH4Cl
D) 1.0 mole C5H5N and 1.5 mole C5H5NHCl
E) none of these

73. The electron configuration for the barium atom is:
A) [Xe] 6s2
B) 1s22s22p63s23p63d104s2
C) 1s22s22p63s23p64s2
D) 1s22s22p63s23p64s1
E) none of these

74. In which reaction is DS° expected to be positive?
A) I2(g) ------à I2(s)
B) 2O2(g) + 2SO(g) ------à 2SO3(g)
C) CH3OH(g) + (3/2)O2(g) ------à CO2(g) + 2H2O(l)
D) H2O(1) ------àH2O(s)
E) none of these

75. At 0°C, the ion-product constant of water, Kw, is 1.2 ´ 10–15. The pH of pure water at 0°C is:
A) 7.46
B) 6.88
C) 7.56
D) 7.00
E) none of these

76. Which part of an animal cell contains the chromosomes?
A) the mitochondrion
B) the nucleus
C) the lysosomes
D) the ribosomes
E) the cytoplasm

77. Indicate the mass action expression for the following reaction:

C) [X]2[Y][W]3[V]

78. Which of the following statements is (are) true?
I. An excited atom can return to its ground state by absorbing
electromagnetic radiation.
II. The energy of an atom is increased when electromagnetic radiation is
emitted from it.
III. The energy of electromagnetic radiation increases as its frequency
IV. An electron in the n = 4 state in the hydrogen atom can go to the
n = 2 state by emitting electromagnetic radiation at the appropriate
V. The frequency and wavelength of electromagnetic radiation are inversely
proportional to each other.
B) I, II, IV

79. Which statement below is not upheld by the second law of thermodynamics?
A) Machines always waste some energy.
B) A machine is never 100% efficient.
C) The entropy of a perfect crystal at 0 K is zero.
D) The change of entropy of the universe is always positive.
E) All of these

80. Name the following:

A) methylbutane
B) isopropane
C) n-pentane
D) dodecane
E) methylpentane

81. The following data were obtained for the reaction of NO with O2. Concentrations are in molecules/cm3 and rates are in molecules/cm3 × s.
[NO]0 [O2]0 Initial Rate
1 ´ 1018 1 ´ 1018 2.0 ´ 1016
2 ´ 1018 1 ´ 1018 8.0 ´ 1016
3 ´ 1018 1 ´ 1018 18.0 ´ 1016
1 ´ 1018 2 ´ 1018 4.0 ´ 1016
1 ´ 1018 3 ´ 1018 6.0 ´ 1016
Which of the following is the correct rate law?
A) Rate = k[NO]2
B) Rate = k[NO][O2]
C) Rate = k[NO]2[O2]2
D) Rate = k[NO][O2]2
E) Rate = k[NO]2[O2]

82. A strip of copper is placed in a 1 M solution of copper nitrate and a strip of silver is placed in a 1 M solution of silver nitrate. The two metal strips are connected to a voltmeter by wires and a salt bridge connects the solutions. The following standard reduction potentials apply:
Ag+(aq) + e– ------à Ag(s) E° = +0.80 V
Cu2+(aq) + 2e– ------à Cu(s) E° = +0.34 V
When the voltmeter is removed and the two electrodes are connected by a wire, which of the following does not take place?
A) Negative ions pass through the salt bridge from the silver half-cell to the copper half-cell.
B) There is a net general movement of silver ions through the salt bridge to the copper half-cell.
C) The silver electrode increases in mass as the cell operates.
D) Some positive copper ions pass through the salt bridge from the copper half-cell to the silver half-cell.
E) Electrons flow in the external circuit from the copper electrode to the silver electrode.

83. Which of the following shows a decrease in entropy?
A) precipitation
B) melting ice
C) two of these
D) a burning piece of wood
E) gaseous reactants forming a liquid

84. In deciding which of two acids is the stronger, one must know:
A) the concentration of each acid solution
B) the pH of each acid solution
C) the equilibrium constant of each acid
D) all of the above
E) both a and c must be known

85. Name the following:

A) propane
B) butane
C) pentane
D) hexane
E) ethane

86. Atoms having greatly differing electronegativities are expected to form:
A) covalent bonds
B) no bonds
C) ionic bonds
D) nonpolar covalent bonds
E) polar covalent bonds

87. The following reaction occurs in basic solution:
F2 + H2O ------à O2 + F–
When the equation is balanced, the sum of the coefficients is:
A) 10
B) 13
C) 12
D) 11
E) none of these

88. Which of the following statements about liquids is true?
A) Droplet formation occurs because of the higher stability associated with increased surface area.
B) The boiling point of a solution is dependent solely on the atmospheric pressure over the solution.
C) Substances that can form hydrogen bonds will display lower melting points than predicted from periodic trends.
D) Liquid rise within a capillary tube because of the small size lowers the effective atmospheric pressure over the surface of the liquid.
E) London dispersion forces arise from a distortion of the electron clouds within a molecule or atom.

89. For the stepwise dissociation of aqueous H3PO4, which of the following is not a conjugate acid-base pair?
A) HPO42– and PO43–
B) H2PO4– and PO43–
C) H3O+ and H2O
D) H3PO4 and H2PO4–
E) H2PO4– and HPO42–

90. Which one of the following decreases as the strength of the attractive intermolecular forces increases?
A) The normal boiling temperature.
B) The extent of deviations from the ideal gas law.
C) The vapor pressure of a liquid.
D) The heat of vaporization.
E) The sublimation temperature of a solid.

91. The number of orbitals having a given value of l is equal to
A) the number of lobes in each orbital
B) 2n + 2
C) 3l
D) 2l + 1
E) l + ml

92. If the change in entropy of the surroundings for a process at 451 K and constant pressure is –326 J/K, what is the heat flow absorbed by for the system?
A) 24.2 kJ
B) 147 kJ
C) 326 kJ
D) 12.1 kJ
E) –147 kJ

93. Identify the missing particle in the following equation:

E) none of these

94. Given the following information:
Li(s) ------à Li(g) heat of sublimation of Li(s) = 166 kJ/mol
HCl(g) ------à H(g) + Cl(g) bond energy of HCl = 427 kJ/mol
Li(g) ------à Li+(g) + e– ionization energy of Li(g) = 520. kJ/mol
Cl(g) + e– ------à Cl–(g) electron affinity of Cl(g) = –349 kJ/mol
Li+(g) + Cl–(g) ------à LiCl(s) lattice energy of LiCl(s) = –829 kJ/mol
H2(g) ------à 2H(g) bond energy of H2 = 432 kJ/mol
calculate the net change in energy for the reaction 2Li(s) + 2HCl(g) ------à 2LiCl(s) + H2(g)
A) –562 kJ
B) –73 kJ
C) 363 kJ
D) –179 kJ
E) None of these


Which of the following result(s) in an increase in the entropy of the system?
I. (See diagram above.)
II. Br2(g) ------à Br2(l)
III. NaBr(s) ------à Na+(aq) + Br-(aq)
IV. O2(298 K) ------à O2(373 K)
V. NH3(1 atm, 298 K) ------à NH3(3 atm, 298 K)
A) I, II, III, V
C) I
E) II, V

96. Atoms which are sp2 hybridized form ____ pi bond(s).
A) 3
B) 0
C) 4
D) 2
E) 1

97. The hybridization of the phosphorus atom in the cation PH2+ is:
A) sp3
B) sp2
C) sp
D) dsp
E) none of these

98. Which of the following is the best description of a protein?
A) two antiparallel chains of nucleic acids connected by hydrogen bonding
B) an alternating chain of amino acids and nucleic acids
C) a chain of amino acids formed by condensation polymerization
D) a chain of nucleotides connected by phosphodiester bonds
E) a chain of amino acids connected by ester bonds

Use the following to answer question 99:

Consider the following portion of the energy-level diagram for hydrogen:
n = 4 –0.1361 ´ 10–18 J
n = 3 –0.2420 ´ 10–18 J
n = 2 –0.5445 ´ 10–18 J
n = 1 –2.178 ´ 10–18 J

99. For which of the following transitions does the light emitted have the longest wavelength?
A) n = 4 to n = 3
B) n = 2 to n = 1
C) n = 3 to n = 2
D) n = 4 to n = 1
E) n = 4 to n = 2

100. Radioactive elements decay via first-order kinetics. Consider a certain type of nucleus that has a rate constant of 1.0 ´ 10–3 h–1. A sample contains 5.0 ´ 109 radioactive nuclides. Calculate the number of nuclides remaining after 39 days have passed.
A) 7.8 ´ 10–11
B) 5.0 ´ 109
C) 64
D) 2.0 ´ 109
E) 2.5 ´ 109

101. A 100-mL sample of water is placed in a coffee cup calorimeter. When 1.0 g of an ionic solid is added, the temperature decreases from 21.5°C to 20.8°C as the solid dissolves. For the dissolving of the solid
A) DSsys< 0
B) DH > 0
C) DSuniv > 0
D) DSsurr > 0
E) none of these

Use the following to answer question 102:

Consider the following data concerning the equation:
H2O2 + 3I– + 2H+ ------à I3– + 2H2O
[H2O2] [I–] [H+] rate
I. 0.100 M 5.00 × 10–4 M 1.00 × 10–2 M 0.137 M/sec
II. 0.100 M 1.00 × 10–3 M 1.00 × 10–2 M 0.268 M/sec
III. 0.200 M 1.00 × 10–3 M 1.00 × 10–2 M 0.542 M/sec
IV. 0.400 M 1.00 × 10–3 M 2.00 × 10–2 M 1.084 M/sec

102. The rate law for this reaction is
A) rate = k[H2O2][H+]
B) rate = k[H2O2][I–]
C) rate = k[I–][H+]
D) rate = k[H2O2]2[I–]2[H+]2
E) rate = k[H2O2][I–][H+]

103. The balanced equation for the reaction of bromate ion with bromide in acidic solution is given by:

At a particular instant in time, the value of –D[Br–]/Dt is 2.0 ´ 10–3 mol/L × s. What is the value of D[Br2]/Dt in the same units?
A) 6.0 ´ 10–3
B) 3.3 ´ 10–5
C) 1.2 ´ 10–3
D) 2.0 ´ 10–3
E) 3.3 ´ 10–3

104. Which of the following indicates the most acidic solution?
A) [OH–] = 0.5 M
B) [H+] = 1 ´ 10–4 M
C) pH = 1.2
D) pOH = 5.9
E) [H+] = 0.3 M

105. The second law of thermodynamics states that
A) the entropy of the universe is constant.
B) the energy of the universe is increasing.
C) the energy of the universe is constant.
D) the entropy of a perfect crystal is zero at 0 K.
E) the entropy of the universe is increasing.

106. The statement that "the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals" is known as
A) Hund's rule.
B) Heisenberg uncertainty principle.
C) the quantum model.
D) the aufbau principle.
E) the Pauli exclusion principle.

107. If, at a given temperature, the equilibrium constant for the reaction

is Kp, then the equilibrium constant for the reaction

can be represented as:
A) Kp2

108. Radioactive elements decay via first-order kinetics. Consider a certain type of nucleus that has a rate constant of 1.0 ´ 10–3 h–1. A sample contains 5.0 ´ 109 radioactive nuclides. Calculate the time required to reduce that number to 2.5 ´ 109.
A) 1.0 ´ 103 h
B) 5.0 ´ 102 h
C) 6.9 ´ 102 h
D) 2.0 ´ 103 h
E) 6.9 ´ 10–4 h

109. Which one of the following statements about the structure of proteins is incorrect?
A) Heat can disrupt tertiary structure.
B) Disulfide bonds provide strong intrachain interactions.
C) Ionized amino acid side chains can form salt bridges within a protein.
D) Hydrogen bonding stabilizes the a-helix proteins.
E) Nonpolar groups tend to face the outside of a protein in an aqueous solution.

110. The ratio of the atomic radius to the nuclear radius is approximately:
A) 105
B) 10–15
C) 102
D) 10–5
E) 1015

111. Which of the following is true for the cell shown here?
Zn(s) ú Zn2+(aq) ------àCr3+(aq) ú Cr(s)
A) The chromium is oxidized.
B) The electrons flow from the cathode to the anode.
C) The zinc is reduced.
D) The electrons flow from the chromium to the zinc.
E) The electrons flow from the zinc to the chromium.

112. Which of the following chemical or physical changes is an endothermic process?
A) the freezing of water
B) the mixing of sulfuric acid and water
C) the evaporation of water
D) the combustion of gasoline
E) none of these

113. Which statement is true?
A) All real processes are irreversible.
B) All statements are true.
C) In a reversible process, the state functions of the system are always much greater than those of the surroundings.
D) There is always more heat given off to the surroundings in a reversible process than in an unharnessed one.
E) A thermodynamically reversible process takes place infinitely fast.

114. Calculate the molality of C2H5OH in a water solution that is prepared by mixing 50.0 mL of C2H5OH with 100.0 mL of H2O at 20°C. The density of the C2H5OH is 0.789 g/mL at 20°C.
A) 0.086 m
B) 1.24 m
C) 0.094 m
D) 8.56 m
E) none of these

115. Which of the following statements is true?
A) The concentration of the products equals that of reactants and is constant at equilibrium.
B) Catalysts are an effective means of changing the position of an equilibrium.
C) When two opposing processes are proceeding at identical rates, the system is at equilibrium.
D) An endothermic reaction shifts toward reactants when heat is added to the reaction.
E) None of these statements is true.

116. The nuclide is radioactive. When one of these atoms decays, a series of a and b-particle emissions occurs, taking the atom through many transformations to end up as an atom of . How many a particles are emitted in converting into ?
A) 2
B) 214
C) 6
D) 4
E) 8

117. Which of the following pairs is isoelectronic?
A) Li+ and K+
B) Na+ and Ne
C) S2– and Ne
D) Al3+ and B3+
E) I– and Cl–

118. Order the intermolecular forces (dipole-dipole, London Dispersion, ionic, and hydrogen-bonding) from weakest to strongest.
A) London Dispersion, ionic, dipole-dipole, and hydrogen-bonding
B) dipole-dipole, ionic, London Dispersion, and hydrogen-bonding
C) London Dispersion, dipole-dipole, hydrogen-bonding, ionic
D) dipole-dipole, London Dispersion, ionic, and hydrogen-bonding
E) hydrogen-bonding, dipole-dipole, London Dispersion, and ionic

119. Which of the following is incorrect?
A) The importance of the equation E = mc2 is that energy has mass.
B) Electromagnetic radiation can be thought of as a stream of particles called photons.
C) All of these are correct.
D) The energy of matter is not continuous and is actually quantized.
E) Energy can only occur in discrete units called quanta.

120. Which of the following statements concerning equilibrium is not true?
A) A system that is disturbed from an equilibrium condition responds in a manner to restore equilibrium.
B) Equilibrium in molecular systems is dynamic, with two opposing processes balancing one another.
C) The value of the equilibrium constant for a given reaction mixture is the same regardless of the direction from which equilibrium is attained.
D) A system moves spontaneously toward a state of equilibrium.
E) The equilibrium constant is independent of temperature.

121. Consider the following system at equilibrium:

Which of the following changes will shift the equilibrium to the right?
I. increasing the temperature
II. decreasing the temperature
III. increasing the volume
IV. decreasing the volume
V. removing some NH3
VI. adding some NH3
VII. removing some N2
VIII. adding some N2

122. A first-order reaction is 35% complete at the end of 55 minutes. What is the value of the rate constant?
A) 7.8 × 10–3 min–1
B) 89 min–1
C) 36 min–1
D) 1.9 × 10–3 min-1
E) none of these

123. Which one of the following types of radiation has the shortest wavelength, the greatest energy, and the highest frequency?
A) none because short wavelength is associated with low energy and low frequency, not high energy and high frequency
B) visible red light
C) infrared radiation
D) ultraviolet radiation
E) visible blue light

124. How many milliliters of 18.4 M H2SO4 are needed to prepare 600.0 mL of 0.10 M H2SO4?
A) 4.0 mL
B) 4.6 mL
C) 1.8 mL
D) 2.7 mL
E) 3.3 mL

125. Which of the following statements is incorrect?
A) Linear molecules cannot have a net dipole moment.
B) A molecule with very polar bonds can be nonpolar.
C) Ionic bonding results from the transfer of electrons from one atom to another.
D) The electrons in a polar bond are found nearer to the more electronegative element.
E) Dipole moments result from the unequal distribution of electrons in a molecule.

126. Which of the following is a conjugate acid/base pair?
A) H3O+/OH–
B) NH4+/NH3
C) HCl/OCl–
D) H2SO4/SO42–
E) none of these

127. How many electrons are transferred in the following reaction?
2ClO3– + 12H+ + 10I– ------à 5I2 + Cl2 + 6H2O
A) 10
B) 12
C) 5
D) 30
E) 2

128. Calculate the [H+] in a solution that has a pH of 11.70.
A) 11.7 M
B) 5.0 ´ 10–3 M
C) 2.0 ´ 10–12 M
D) 2.3 M
E) none of these

129. How many of the following molecules possess dipole moments?
BH3, CH4, PCl5, H2O, HF, H2
A) 4
B) 5
C) 3
D) 1
E) 2

130. A correct statement of Henry's law is:
A) the concentration of a gas in solution is independent of pressure.
B) the concentration of a gas in solution is directly proportional to the mole fraction of solvent.
C) the concentration of a gas in a solution is inversely proportional to pressure.
D) the concentration of a gas in solution is inversely proportional to temperature.
E) none of these

131. Which of the following bonds is least polar?
A) C–O
B) S–Cl
C) They are all nonpolar.
D) H–C
E) Br–Br

132. The average rate of disappearance of ozone in the reaction 2O3(g) ------à 3O2(g) is found to be 9.0 ´ 10–3 atm over a certain interval of time. What is the rate of appearance of O2 during this interval?
A) 2.7 ´ 10–5 atm/s
B) 1.3 ´ 10–2 atm/s
C) 9.0 ´ 10–3 atm/s
D) 6.0 ´ 10–3 atm/s
E) 3.0 ´ 10–5 atm/s

133. The lattice energy of NaI is 686 kJ/mol and its heat of solution is –7.6 kJ/mol. Calculate the hydration of energy of NaI(s).
A) –678
B) +694
C) +678
D) +15.2
E) –694

134. A two-bulbed flask contains seven particles. What is the probability of finding all seven particles on the left side?
A) 0.13%
B) 0.32%
C) 0.78%
D) 3.1%
E) 0.93%

135. Using the following bond energies
Bond Bond Energy (kJ/mol)
CºC 839
C–H 413
O=O 495
C=O 799
O–H 467
estimate the heat of combustion for one mole of acetylene:
C2H2(g) + (5/2)O2(g) ------à 2CO2(g) + H2O(g)
A) +365 kJ
B) +447 kJ
C) –447 kJ
D) –1228 kJ
E) 1228 kJ

136. Knowing that DHvap for water is 40.7 kJ/mol, calculate Pvap of water at 37°C.
A) 6.90 torr
B) 25.4 torr
C) 52.6 torr
D) 18.7 torr
E) 12.4 torr

137. Consider the following reaction:
The equilibrium constant K is 0.28 at 900°C. What is Kp at this temperature?
A) 4.0 ´ 10–5
B) 1.0 ´ 10–5
C) 5.0 ´ 10–5
D) 3.0 ´ 10–5
E) 2.0 ´ 10–5

138. In the following fission reaction, identify the other product:


139. Radioactive tracers are useful in studying very low concentrations of chemical species. A chemist has a sample of HgI2 in which part of the iodine is the radioactive nuclide of mass 131, so that the count rate is 5.0 ´ 1011 counts per minute per mole of I. The solid mercuric iodide is placed in water and allowed to come to equilibrium. Then 100 mL of the solution is withdrawn, and its radioactivity is measured and found to give 22 counts per minute. What is the molar concentration of iodide ion in the solution?
A) 4.4 ´ 10–10
B) 4.4 ´ 10–11
C) 1.1 ´ 10–9
D) 1.1 ´ 10–10
E) 1.1 ´ 10–11

140. A 20.0-g sample of methyl alcohol (CH3OH, molar mass = 32.0 g/mol) was dissolved in 30.0 g of water. The mole fraction of CH3OH is:
A) 0.728
B) 0.400
C) 0.625
D) 0.667
E) none of these

141. The half-life for electron capture for is 1.3 billion years. What will be the ratio in a rock that is 4.5 billion years old?
A) 11.
B) 0.10
C) 10.
D) 0.36
E) 0.091

142. A weak acid, HF, is in solution with dissolved sodium fluoride, NaF. If HCl is added, which ion will react with the extra hydrogen ions from the HCl to keep the pH from changing?
A) Na–
B) Na+
C) OH–
D) F–
E) none of these

143. The number of a certain radioactive nuclide present in a sample decays from 1.00 x 103 to 2.50 x 102 in 10 minutes. What is the half-life of this radioactive species?
A) not enough information given
B) 20 minutes
C) 10 minutes
D) 2.5 minutes
E) 5 minutes

144. An element with the electron configuration [Xe]4f145d76s2 would belong to which class on the periodic table?
A) transition elements
B) halogens
C) alkaline earth elements
D) rare earth elements
E) none of these

145. A sample of wood from an Egyptian mummy case gives a 14C count of 9.4 cpm/gC (counts per minute per gram of carbon). How old is the wood? (The initial decay rate of 14C is 15.3 cpm/gC, and its half-life is 5730 years.)
A) 4570 yr
B) 4030 yr
C) 3420 yr
D) 6400 yr
E) none of these

146. In chemical reactions involving alkenes and alkynes, pi bonds can be broken and replaced with sigma bonds. These reactions are called
A) addition reactions.
B) polymerization reactions.
C) combustion reactions.
D) dehydrogenation reactions.
E) substitution reactions.

147. Which of the following atoms or ions has 3 unpaired electrons?
A) O
B) N
C) S2–
D) Zn2+
E) Al

148. In the gaseous phase, which of the following diatomic molecules would be the most polar?
A) NaCl
B) NaF
C) LiF
D) CsF
E) CsCl

149. If two pyramid-shaped dice (with numbers 1 through 4 on the sides) were tossed, which outcome has the highest entropy?
A) The sum of the dice is 6.
B) The sum of the dice is 7.
C) The sum of the dice is 5.
D) The sum of the dice is 4.
E) The sum of the dice is 3.

150. Hydrogen bonds account for which of the following observation?
A) Water molecules are bent or "V-shaped".
B) For its molar mass, water has a high boiling point.
C) Air is more dense than hydrogen gas.
D) Hydrogen is easily combustible with oxygen.
E) Hydrogen naturally exists as a diatomic molecule.

151. Which of these statements about benzene is true?
A) Benzene is an example of a molecule that displays ionic bonding.
B) Benzene contains only p bonds between C atoms.
C) All carbon atoms in benzene are sp3 hybridized.
D) The bond order of each C–C bond in benzene is 1.5.
E) All of these statements are false.

152. In a 0.1 molar solution of NaCl in water, which one of the following will be closest to 0.1?
A) The molality of NaCl
B) the mass percent of NaCl
C) the mass fraction of NaCl
D) the mole fraction of NaCl
E) all of these are about 0.1

153. How many electrons are transferred in the following reaction?
Fe + 2HCl ------à FeCl2 + H2
A) 4
B) not enough information given
C) 0
D) 2
E) 1

154. Which of the following reactions is possible at the anode of a galvanic cell?
A) two of these
B) Zn + Cu2+ ------à Zn2+ + Cu
C) Zn2+ + Cu ------à Zn + Cu2+
D) Zn2+ + 2e– ------à Zn
E) Zn ------à Zn2+ + 2e–

155. Which of the following is not determined by the principal quantum number, n, of the electron in a hydrogen atom?
A) the minimum wavelength of the light needed to remove the electron from the atom.
B) the size of the corresponding atomic orbital(s)
C) the shape of the corresponding atomic orbital(s)
D) the energy of the electron
E) All of these are determined by n.

156. Which of the following species is best described by drawing resonance structures?
A) PH3
C) NH4+
D) SO3–
E) O3

157. Which of the following is a reasonable criticism of the Bohr model of the atom?
A) It makes no attempt to explain why the negative electron does not eventually fall into the positive nucleus.
B) It does not adequately predict the ionization energy of the 1st energy level electrons for one-electron species for elements other than hydrogen.
C) It shows the electrons to exist outside of the nucleus.
D) It does not adequately predict the line spectrum of hydrogen.
E) It does not adequately predict the ionization energy of the valence electron(s) for elements other than hydrogen.

158. How many isomers of C4H8 are there?
A) 1
B) 2
C) 6
D) 5
E) 3

159. 2.5 L of an aqueous solution containing 50.00 g of MgCl2 dissolved in pure water is prepared. The molarity of the solution is:
A) 1.1 M
B) 0.59 M
C) 0.47 M
D) 0.82 M
E) 0.21 M

160. For which process is DS negative?
A) grinding a large crystal of KCl to powder
B) compressing 1 mol Ne at constant temperature from 1.5 atm to 0.5 atm
C) evaporation of 1 mol of CCl4(l)
D) mixing 5 mL ethanol with 25 mL water
E) raising the temperature of 100 g Cu from 275 K to 295 K

Friday, March 15, 2002

Do all work on a separate paper. Show all work. This is due WED March 20, 2002.
Go to the following website:

1. Go to the section on Acids/Bases

1. Discuss the five properties of acids and bases listed in the internet article.

2. What is “aqua regia”? What is this known for?

3. What was Antoine Lavoisier’ s basic hypothesis concerning the cause of acidity? Is this idea still considered valid today?

4. How did Joseph Priestley discover HCl? Write balanced chemical equations to show the reactions that lead to this discovery.

5. Discuss Humphry Davy’s contribution to acid/base theory.

6. What was Justus Liebig’s definition of an acid?

7. What was Svante Arrhenius’ most important contribution to chemistry?

8. Define an acid and a base according to Arrhenius.

9. Discuss the fundamental problem with the Arrhenius definition of an acid.

10. Define an acid and a base according to Brønsted and Lowry. Discuss the problem with calling an acid a proton “donor”.

11. Define a conjugate pair (according to Brønsted and Lowry).
12. Identify the conjugate acid base pairs:
HC2H3O2 + H2O <===> H3O+ + C2H3O2¯

NH3 + H2O <===> NH4+ + OH¯
12. Define pH.

13. Discuss the relationship between significant figures and the pH scale.

14. Do the 6 pH Practice Problems. Show all work!

15. Do the 6 pOH Practice Problems. Show all work!

Tuesday, February 26, 2002

Honors Chemistry Equilibrium Internet Assignment (100 points)

Go to the following website:

1. Go to the section on Equilibrium.

a. What determines the exact moment of equilibrium in a chemical reaction?
b. At equilibrium what can you say about the concentration of the reactants and the products?
c. Do all chemical reactions reach equilibrium? Explain.
2. Go to the section “Metaphors………” Click on “my description ……..” Do parts (a) and (b).

3. Click on “The Equilibrium Constant” Define the Law of Mass Action. Do the 5 practice problems. Show your work. Check your answers.

4 . At equilibrium, what is the relationship between reaction rates of the forward and reverse reactions?

5. Go to: “Calculating the equilibrium constant……” Work through the examples. Show your work.

6. Go on to the next section, “Calculating Equilibrium Concentrations from Initial Conditions. Work through the examples. Show your work.

7. Go on to “lets do some more……..” Note: These solutions require using the quadratic equation. Show your work.

8. Go on to “LeChatlier’s Principle” Express this principle in your own words.

9. What are three factors that affect a chemical reaction at equilibrium? Work through the examples, note: the answers are given. Show your work.

10.Go to: “more LeChatlier problems “ Work these problems. Show your work. Check your answers.

Monday, February 04, 2002

1. Go to website:
2. Click on “Bonding & Isomerism”

3. The Table of Contents includes 13 sections.

4. Read and take notes on each section.

5. Answer the following questions using complete sentences. (5 points each) 85 points total.

1. What fundamental force is responsible for chemical bonding?
2. J.J. Thompson's bonding theory (1904) had two problems. What were they, and how were they solved?
3. What is the central idea of an ionic bond?
4. In the section: "Which Element form Ionic Bonds" what are the three main points?
5. Who was the first to quantify electonegativity?
6. What are the three principle bond types?
7. What are the electronegativity differences associated with the three bond types?
8. If the electronegativity difference is between 1.6 and 2.0 what is the rule for bond type?
9. Read the rules for drawing Lewis structures and work through the examples given. Then draw the Lewis structure for PF3 and PF5.
10. Read the rules for Expanded and Deficient Octets and work through the examples. Then draw the Lewis structures for SF6 and BrF5
11. Read the section on VSEPR structures of odd electron molocules. How does the bond angle on ClO+2 change from the bond angle on ClO2.
12. What is the general principle in predicting molecular polarity?
13. Compare and contrast bond polarity with molecular polarity.
14. What is formal charge?
15. Explain how to determine formal charge.
16. Define resonance.
17. Determine the resonance structures for ozone, carbon monoxide and carbon dioxide.
18. 15 points. EXTRA CREDIT: Do some research and find the five resonance structures for benzene.

Wednesday, January 23, 2002

Electrons in Atoms and Spectroscopy Internet Assignment

1. Go to website:

2. Click on “Electrons in Atoms & Spectroscopy”

3. The Table of Contents includes 4 sections.

4. Read and take notes on each section.

5. Answer the following questions using complete sentences.

a. Discuss the relationship between emission and absorption spectra.
b. Discuss the relationship between line and continuous spectra.
c. Explain the “Balmer series”
d. What is a wave number? Derive an equation for the energy of a photon as a function of wave number.
e. Summarize the “assumptions” of the Bohr model.
f. List the four features of the Bohr model.

Tuesday, January 08, 2002

I have posted 150 sample questions for the final exam (with answers). I recommend you work through each of these problems and try to understand concepts. Do NOT just memorize answers. The final exam will cover the same concepts and the questions will be in the same format. The final exam will contain 75 questions covering chapters 1-6. Study hard. Good luck.
1. According to the law of definite proportions:
A) if the same two elements form two different compounds, they do so in the same ratio.
B) it is not possible for the same two elements to form more than one compound.
C) the ratio of the masses of the elements in a compound is always the same.
D) the total mass after a chemical change is the same as before the change.

2. Many classic experiments have given us indirect evidence of the nature of the atom. Which of the experiments listed below did not give the results described?
A) The Rutherford experiment proved the Thomson "plum-pudding" model of the atom to be essentially correct.
B) The Rutherford experiment was useful in determining the nuclear charge on the atom.
C) Millikan's oil-drop experiment showed that the charge on any particle was a simple multiple of the charge on the electron.
D) The electric discharge tube proved that electrons have a negative charge.

3. Which of the following pairs can be used to illustrate the law of multiple proportions?
A) SO and SO2
B) CO and CaCo3
C) H2O and C12H22O11
D) H2SO4 and H2S
E) KCl and KClO2

4. The first scientist to show that atoms emit any negative particles was
A) J. J. Thomson.
B) Lord Kelvin.
C) Ernest Rutherford.
D) William Thomson.
E) John Dalton.

5. The scientist whose alpha-particle scattering experiment led him to conclude that the nucleus of an atom contains a dense center of positive charge is
A) J. J. Thomson.
B) Lord Kelvin.
C) Ernest Rutherford.
D) William Thomson.
E) John Dalton.

6. Which one of the following statements about atomic structure is false?
A) The electrons occupy a very large volume compared to the nucleus.
B) Almost all of the mass of the atom is concentrated in the nucleus.
C) The protons and neutrons in the nucleus are very tightly packed.
D) The number of protons and neutrons is always the same in the neutral atom.

7. Which of the following name(s) is(are) correct?
1. sulfide S2–
2. ammonium chloride NH4Cl
3. acetic acid HC2H3O2
4. barium oxide BaO
A) all
B) none
C) 1, 2
D) 3, 4
E) 1, 3, 4

8. Which of the following atomic symbols is incorrect?

9. The element rhenium (Re) exists as two stable isotopes and 18 unstable isotopes. Rhenium-185 has in its nucleus
A) 75 protons, 75 neutrons.
B) 75 protons, 130 neutrons.
C) 130 protons, 75 neutrons.
D) 75 protons, 110 neutrons.
E) not enough information is given.

10. Which statement is not correct?
A) The mass of an alpha particle is 7300 times that of the electron.
B) An alpha particle has a 2+ charge.
C) Three types of radioactive emission are gamma rays, beta rays, and alpha particles.
D) A gamma ray is high-energy "light."
E) There are only three types of radioactivity known to scientists today.

11. All of the following are true except:
A) Ions are formed by adding electrons to a neutral atom.
B) Ions are formed by changing the number of protons in an atom's nucleus.
C) Ions are formed by removing electrons from a neutral atom.
D) An ion has a positive or negative charge.
E) Metals tend to form positive ions.

12. Which among the following represent a set of isotopes? Atomic nuclei containing:
I. 20 protons and 20 neutrons.
II. 21 protons and 19 neutrons.
III. 22 neutrons and 18 protons.
IV. 20 protons and 22 neutrons.
V. 21 protons and 20 neutrons.
C) I, V
D) I, IV and II, V
E) No isotopes are indicated.

13. The average mass of a carbon atom is 12.011. Assuming you were able to pick up only one carbon unit, the chances that you would randomly get one with a mass of 12.011 is
A) 0%.
B) 0.011%.
C) about 12%.
D) 12.011%.
E) greater than 50%.

14. An ion is formed
A) by either adding or subtracting protons from the atom.
B) by either adding or subtracting electrons from the atom.
C) by either adding or subtracting neutrons from the atom.
D) All of the above are true.
E) Two of the above are true.

15. How many oxygen atoms are there in one formula unit of Ca3(PO4)2?
A) 2
B) 4
C) 6
D) 8
E) none of these

16. All of the following are characteristics of metals except:
A) good conductors of heat
B) malleable
C) ductile
D) often lustrous
E) tend to gain electrons in chemical reactions

17. A species with 12 protons and 10 electrons is
A) Ne2+
B) Ti2+
C) Mg2+
D) Mg
E) Ne2–

18. The correct name for LiCl is
A) lithium monochloride
B) lithium (I) chloride
C) monolithium chloride
D) lithium chloride
E) monolithium monochloride

19. The correct name for FeO is
A) iron oxide
B) iron (II) oxide
C) iron (III) oxide
D) iron monoxide
E) iron (I) oxide

20. The formula for calcium bisulfate is
A) Ca(SO4)2
B) CaS2
C) Ca(HSO4)2
D) Ca2HSO4
E) Ca2S

21. Which of the following is incorrectly named?
A) Pb(NO3)2, lead(II) nitrate
B) NH4ClO4, ammonium perchlorate
C) PO43–, phosphate ion
D) Mg(OH)2, magnesium hydroxide
E) NO3–, nitrite ion

22. All of the following are in aqueous solution. Which is incorrectly named?
A) H2SO4, sulfuric acid
B) H2CO3, carbonic acid
C) H3PO4, phosphoric acid
D) HCN, cyanic acid
E) HCl, hydrochloric acid

23. It is estimated that uranium is relatively common in the earth's crust, occurring in amounts of 4 g/metric ton. A metric ton is 1000 kg. At this concentration, what mass of uranium is present in 1.0 mg of the earth's crust?
A) 4 nanograms
B) 4 micrograms
C) 4 milligrams
D) 4 ´ 10–5 g
E) 4 centigrams

24. Which of the following is an example of a qualitative observation?
A) A piece of wood is 5.3 cm long.
B) Solution 1 is much darker than solution 2.
C) The volume of liquid in beaker A is 4.3 mL
D) The temperature of the liquid is 60°C.
E) none of these

25. A quantitative observation
A) contains a number and a unit.
B) does not contain a number.
C) always makes a comparison.
D) must be obtained through experimentation.
E) none of these

26. A set of tested hypotheses that gives an overall explanation of some natural phenomenon is called a(n)
A) observation.
B) measurement.
C) theory.
D) natural law.
E) experiment.

27. Express 784000000 in exponential notation.
A) 7.84 ´ 106
B) 7.84 ´ 108
C) 78.4 ´ 107
D) 784 ´ 106
E) 784 ´ 107

28. A titration was performed to find the concentration of hydrochloric acid with the following results:
Trial Molarity
1 1.25 ± 0.01
2 1.24 ± 0.01
3 1.26 ± 0.01
The actual concentration of HCl was determined to be 1.000 M; the results of the titration are:
A) both accurate and precise.
B) accurate but imprecise.
C) precise but inaccurate.
D) both inaccurate and imprecise.
E) accuracy and precision are impossible to determine with the available information.

29. Which of the following is the least probable concerning five measurements taken in the lab?
A) The measurements are accurate and precise.
B) The measurements are accurate but not precise.
C) The measurements are precise but not accurate.
D) The measurements are neither accurate nor precise.
E) All of these are equally probable.

30. The amount of uncertainty in a measured quantity is determined by:
A) both the skill of the observer and the limitations of the measuring instrument.
B) neither the skill of the observer nor the limitations of the measuring instrument.
C) the limitations of the measuring instrument only.
D) the skill of the observer only.

31. A scientist obtains the number 1250.37986 on a calculator. If this number actually has four (4) significant figures, how should it be written?
A) 1251
B) 1250.3799
C) 1250.4
D) 1.250 ´ 103
E) 1.250 ´ 10–3

32. How many significant figures are there in the number 0.0322?
A) 3
B) 5
C) 4
D) 2
E) 0

33. A piece of indium with a mass of 16.6 g is submerged in 46.3 cm3 of water in a graduated cylinder. The water level increases to 48.6 cm3. The correct value for the density of indium from these data is:
A) 7.217 g/cm3
B) 7.2 g/cm3
C) 0.14 g/cm3
D) 0.138 g/cm3
E) more than 0.1 g/cm3 away from any of these values.

34. Express 0.000543 in exponential notation.
A) 5.43 ´ 10–4
B) 5.43 ´ 10–6
C) 54.3 ´ 10–5
D) 54.3 ´ 10–3
E) 543 ´ 10–3

35. Using the rules of significant figures, calculate the following:

A) 17.5
B) 18
C) 17
D) 20
E) 17.48

36. Express the volume 159 dm3 in liters.
A) 159 L
B) 159000 L
C) 0.159 L
D) 15.9 L
E) 1.59 L

37. The degree of agreement among several measurements of the same quantity is called __________. It reflects the reproducibility of a given type of measurement.
A) accuracy
B) error
C) precision
D) significance
E) certainty

38. Convert 0.092 ft3 to L. (2.54 cm = 1 in., 1 L = 1 dm3)
A) 26 L
B) 2.6 L
C) 3.2 ´ 10–3 L
D) 1.8 L
E) 0.40 L

39. 423 Kelvin equals
A) 150. °F
B) 273. °F
C) 696. °F
D) 150. °C
E) 696. °C

40. The melting point of lead is 327°C. What is this on the Fahrenheit scale?
(TF = TC ´ (9°F/5°C) + 32°F)
A) 620.6°F
B) 600°F
C) 895°F
D) 621°F
E) 547°F

41. The state of matter for an object that has a definite volume but not a definite shape is
A) solid state.
B) liquid state.
C) gaseous state.
D) elemental state.
E) mixed state.

42. In 1928, rhenium cost $10,000/kg. It now costs $40/troy ounce. What is the present cost of a gram of rhenium? (1 troy ounce = 31.10 g)
A) less than $1.00
B) between $1.00 and $10
C) between $10 and $50
D) between $50 and $100
E) over $100

43. Bromine exists naturally as a mixture of bromine-79 and bromine-81 isotopes. An atom of bromine-79 contains
A) 35 protons, 44 neutrons, 35 electrons.
B) 34 protons and 35 electrons, only.
C) 44 protons, 44 electrons, and 35 neutrons.
D) 35 protons, 79 neutrons, and 35 electrons.
E) 79 protons, 79 electrons, and 35 neutrons.

44. The atomic mass of rhenium is 186.2. Given that 37.1% of natural rhenium is rhenium-185, what is the other stable isotope?

45. Gallium consists of two isotopes of masses 68.95 amu and 70.95 amu with abundances of 60.16% and 39.84%, respectively. What is the average atomic mass of gallium?
A) 69.95
B) 70.15
C) 71.95
D) 69.75
E) 69.55

46. Iron is biologically important in the transport of oxygen by red blood cells from the lungs to the various organs of the body. In the blood of an adult human, there are approximately 2.60 ´ 1013 red blood cells with a total of 2.90 g of iron. On the average, how many iron atoms are present in each red blood cell? (molar mass (Fe) = 55.85 g)
A) 8.33 ´ 10–10
B) 1.20 ´ 109
C) 3.12 ´ 1022
D) 2.60 ´ 1013
E) 5.19 ´ 10–2

47. A sample of ammonia has a mass of 56.6 g. How many molecules are in this sample?
A) 3.32 molecules
B) 17.03 ´ 1024 molecules
C) 6.78 ´ 1023 molecules
D) 2.00 ´ 1024 molecules
E) 1.78 ´ 1024 molecules

48. How many moles of hydrogen sulfide are contained in a 35.0-g sample of this gas?
A) 2.16 mol
B) 1.03 mol
C) 7.43 mol
D) 10.4 mol
E) 6.97 mol

49. What is the molar mass of ethanol (C2H5OH)?
A) 45.07
B) 38.90
C) 46.07
D) 34.17
E) 62.07

50. Phosphorus has the molecular formula P4 and sulfur has the molecular formula S8. How many grams of phosphorus contain the same number of molecules as 6.41 g of sulfur?
A) 3.10 g
B) 3.21 g
C) 6.19 g
D) 6.41 g
E) none of these

51. A given sample of xenon fluoride contains molecules of a single type XeFn, where n is some whole number. Given that 9.03 ´ 1020 molecules of XeFn weigh 0.311 g, calculate n.
A) 1
B) 2
C) 4
D) none of these

Use the following to answer question 52:

Phosphoric acid can be prepared by reaction of sulfuric acid with "phosphate rock" according to the equation:
Ca3(PO4)2 + 3H2SO4 ® 3CaSO4 + 2H3PO4

52. Suppose the reaction is carried out starting with 103 g of Ca3(PO4)2 and 75.0 g of H2SO4. Which substance is the limiting reactant?
A) Ca3(PO4)2
B) H2SO4
C) CaSO4
D) H3PO4
E) none of these

53. How many atoms of hydrogen are present in 6.0 g of water?
A) 2.0 × 1023
B) 7.2 × 1024
C) 1.1 × 1024
D) 4.0 × 1023
E) 0.66

54. A substance contains 35.0 g nitrogen, 5.05 g hydrogen, and 60.0 g of oxygen. How many grams of hydrogen are there in a 185-g sample of the substance?
A) 9.34 g
B) 18.7 g
C) 10.6 g
D) 5.05 g
E) 36.6 g

55. Nitric acid contains what percent hydrogen by mass?
A) 20.0%
B) 10.0%
C) 4.50%
D) 1.60%
E) 3.45%

56. In balancing an equation, we change the __________ to make the number of atoms on each side of the equation balance.
A) formulas of compounds in the reactants
B) coefficients of compounds
C) formulas of compounds in the products
D) subscripts of compounds
E) none of these

57. What is the coefficient for water when the following equation is balanced?
As(OH)3(s) + H2SO4(aq) ® As2(SO4)3(aq) + H2O(1)
A) 1
B) 2
C) 4
D) 6
E) 12

58. A compound is composed of element X and hydrogen. Analysis shows the compound to be 80% X by mass, with three times as many hydrogen atoms as X atoms per molecule. Which element is element X?
A) He
B) C
C) F
D) S
E) none of these

59. You heat 3.970 g of a mixture of Fe3O4 and FeO to form 4.195 g Fe2O3. The mass percent of FeO originally in the mixture was:
A) 12.1%
B) 28.7%
C) 71.3%
D) 87.9%
E) none of these

60. You heat 3.970 g of a mixture of Fe3O4 and FeO to form 4.195 g Fe2O3. The mass of oxygen reacted is
A) 0.225 g.
B) 0.475 g.
C) 1.00 g.
D) cannot be determined
E) none of these

61. The limiting reactant in a reaction
A) has the lowest coefficient in a balanced equation.
B) is the reactant for which you have the fewest number of moles.
C) has the lowest ratio of moles available/coefficient in the balanced equation.
D) has the lowest ratio of coefficient in the balanced equation/moles available.
E) none of these

62. Given the equation 3A + B ® C + D, you react 1 mole of A with 3 moles of B. True or false: A is the limiting reactant because you have fewer moles of A than B.

63. Suppose the reaction Ca3(PO4)2 + 3H2SO4 ® 3CaSO4 + 2H3PO4 is carried out starting with 103 g of Ca3(PO4)2 and 75.0 g of H2SO4. How much phosphoric acid will be produced?
A) 74.9 g
B) 50.0 g
C) 112 g
D) 32.5 g
E) 97.6 g

64. A substance, A2B, has the composition by mass of 60% A and 40% B. What is the composition of AB2 by mass?
A) 40% A, 60% B
B) 50% A, 50% B
C) 27% A, 73% B
D) 33% A, 67% B
E) none of these

65. An oxide of iron has the formula Fe3O4. What mass percent of iron does it contain?
A) 0.72%
B) 28%
C) 30.%
D) 70.%
E) 72%

66. A chloride of rhenium contains 63.6% rhenium. What is the formula of this compound?
A) ReCl
B) ReCl3
C) ReCl5
D) ReCl7
E) Re2Cl3

67. A 2.00-g sample of an oxide of bromine is converted to 2.936 g of AgBr. Calculate the empirical formula of the oxide. (molar mass for AgBr = 187.78)
A) BrO3
B) BrO2
C) BrO
D) Br2O
E) none of these

68. Adipic acid contains 49.32% C, 43.84% O, and 6.85% H by mass. What is the empirical formula?
A) C3H5O2
B) C3H3O4
C) C2HO3
D) C2H5O4
E) C3HO3

69. The characteristic odor of pineapple is due to ethyl butyrate, a compound containing carbon, hydrogen, and oxygen. Combustion of 2.78 g of ethyl butyrate leads to formation of 6.32 g of CO2 and 2.58 g of H2O. The properties of the compound suggest that the molar mass should be between 100 and 150. What is the molecular formula?

70. What is the coefficient for oxygen when the following equation is balanced?
NH3(g) + O2(g) ® NO2(g) + H2O(g)
A) 3
B) 6
C) 7
D) 12
E) 14

71. Determine the coefficient for O2 when the following equation is balanced in standard form (smallest whole number integers)
C4H10(g) + O2(g) ® CO2(g) + H2O(g)
A) 4
B) 8
C) 10
D) 13
E) 20

72. The equation, wPCl5 + xH2O ® yPOCl3 + zHCl, is properly balanced when
A) w = 1, x = 2, y = 2, z = 4
B) w = 2, x = 2, y = 2, z = 2
C) w = 2, x = 2, y = 2, z = 1
D) w = 1, x = 1, y = 1, z = 2
E) none of these

73. What would be the g Al/mole S ratio for the product of a reaction between aluminum and sulfur?
A) 26.98 g Al/mol S
B) 53.98 g Al/mol S
C) 80.94 g Al/mol S
D) 17.99 g Al/mol S
E) 40.47 g Al/mol S

74. The Claus reactions, shown below, are used to generate elemental sulfur from hydrogen sulfide.

How much sulfur (in grams) is produced from 48.0 grams of O2?
A) 16.0 g
B) 32.1 g
C) 48.1 g
D) 96.2 g
E) none of these

75. The refining of aluminum from bauxite ore (which contains 50.% Al2O3 by mass) proceeds by the overall reaction 2Al2O3 + 3C ® 4Al + 3CO2. How much bauxite ore is required to give the 5.0 ´ 1013 g of aluminum produced each year in the United States? (Assume 100% conversion.)
A) 1.3 ´ 1013 g
B) 5.3 ´ 1013 g
C) 1.9 ´ 1014 g
D) 7.6 ´ 1014 g

76. SO2 reacts with H2S as follows:
2H2S + SO2 ® 3S + 2H2O
When 7.50 g of H2S reacts with 12.75 g of SO2, which statement applies?
A) 6.38 g of sulfur are formed.
B) 10.6 g of sulfur are formed.
C) 0.0216 moles of H2S remain.
D) 1.13 g of H2S remain.
E) SO2 is the limiting reagent.

77. An unknown substance dissolves readily in water but not in benzene (a nonpolar solvent). Molecules of what type are present in the substance?
A) neither polar nor nonpolar
B) polar
C) either polar or nonpolar
D) nonpolar

78. 1.00 mL of a 3.50 ´ 10–4 M solution of oleic acid is diluted with 9.00 mL of petroleum ether, forming solution A. 2.00 mL of solution A is diluted with 8.00 mL of petroleum ether, forming solution B. How many grams of oleic acid are 5.00 mL of solution B? (molar mass for oleic acid = 282 g/mol)
A) 4.94 ´ 10–4 g
B) 7.00 ´ 10–6 g
C) 4.94 ´ 10–5 g
D) 1.97 ´ 10–6 g
E) 9.87 ´ 10–6 g

79. How many grams of NaCl are contained in 350. mL of a 0.250 M solution of sodium chloride?
A) 41.7 g
B) 5.11 g
C) 14.6 g
D) 87.5 g
E) none of these

80. Which of the following aqueous solutions contains the greatest number of ions?
A) 400.0 mL of 0.10 M NaCl
B) 300.0 mL of 0.10 M CaCl2
C) 200.0 mL of 0.10 M FeCl3
D) 200.0 mL of 0.10 M KBr
E) 800.0 mL of 0.10 M sucrose

81. A 54.8 g sample of SrCl2 is dissolved in 112.5 mL of solution. Calculate the molarity of this solution.
A) 0.346 M
B) 3.07 M
C) 3.96 M
D) 8.89 M
E) none of these

82. A 51.24-g sample of Ba(OH)2 is dissolved in enough water to make 1.20 liters of solution. How many mL of this solution must be diluted with water in order to make 1.00 liter of 0.100 molar Ba(OH)2?
A) 400. mL
B) 333 mL
C) 278 mL
D) 1.20 ´ 103 mL
E) none of these

83. How many grams of NaOH are contained in 5.0 ´ 102 mL of a 0.80 M sodium hydroxide solution?
A) 16 g
B) 80. g
C) 20. g
D) 64 g
E) none of these

84. The net ionic equation for the reaction of aluminum sulfate and sodium hydroxide contains which of the following species?
A) 3Al3+(aq)
B) OH–(aq)
C) 3OH–(aq)
D) 2Al3+(aq)
E) 2Al(OH)3(s)

85. Which of the following is paired incorrectly?
A) HI – strong acid
B) HNO3 – weak acid
C) Ba(OH)2 – strong base
D) HBr – strong acid
E) NH3 – weak acid

86. The interaction between solute particles and water molecules, which tends to cause a salt to fall apart in water, is called
A) hydration.
B) polarization.
C) dispersion.
D) coagulation.
E) conductivity.

87. When solutions of cobalt(II) chloride and carbonic acid react, which of the following terms will be present in the net ionic equation?
A) CoCO3(s)
B) H+(aq)
C) 2CoCO3(s)
D) 2Cl–(aq)
E) two of these

88. You mix 260. mL of 1.20 M lead(II) nitrate with 300. mL of 1.90 M potassium iodide. The lead(II) iodide is insoluble. Which of the following is false?
A) The final concentration of Pb2+ ions is 0.0482 M.
B) You form 131 g of lead(II) iodide.
C) The final concentration of K+ is 1.02 M.
D) The final concentration of NO3– is 1.02 M.
E) All are true.

89. You have 2 solutions of chemical A. To determine which has the highest concentration of A in molarity, what is the minimum number of the following you must know?
I. the mass in grams of A in each solution
II. the molar mass of A
III. the volume of water added to each solution
IV. the total volume of the solution
A) 0
B) 1
C) 2
D) 3
E) You must know all of them.

90. The following reactions
2K(s) + Br2(l) ® 2KBr(s)
AgNO3(aq) + NaCl(aq) ® AgCl(s) + NaNO3(aq)
HCl(aq) + KOH(aq) ® H2O(l) + KCl(aq)
are examples of
A) precipitation reactions.
B) redox, precipitation, and acid-base, respectively.
C) precipitation (two) and acid-base reactions, respectively.
D) redox reactions.
E) none of these

91. Aqueous solutions of sodium sulfide and copper(II) chloride are mixed together. Which statement is correct?
A) Both NaCl and CuS precipitate from solution.
B) No precipitate forms.
C) CuS will precipitate from solution.
D) NaCl will precipitate from solution.
E) No reaction will occur.

92. In the balanced molecular equation for the neutralization of sodium hydroxide with sulfuric acid, the products are:
A) NaSO4 + H2O
B) NaSO3 + 2H2O
C) 2NaSO4 + H2O
D) Na2S + 2H2O
E) Na2SO4 + 2H2O

93. If all of the chloride in a 5.000-g sample of an unknown metal chloride is precipitated as AgCl with 70.90 mL of 0.2010 M AgNO3, what is the percentage of chloride in the sample?
A) 50.55%
B) 10.10%
C) 1.425%
D) 20.22%
E) none of these

94. You have separate solutions of HCl and H2SO4 with the same concentrations in terms of molarity. You wish to neutralize a solution of NaOH. Which acid solution would require more volume (in mL) to neutralize the base?
A) the HCl solution
B) the H2SO4 solution
C) You need to know the acid concentrations to answer this question.
D) You need to know the volume and concentration of the NaOH solution to answer this question.
E) You need to know the acid concentrations and the volume and concentration of the NaOH solution to answer this question.

95. A 3.00-g sample of an alloy (containing only Pb and Sn) was dissolved in nitric acid (HNO3). Sulfuric acid was added to this solution, which precipitated 2.93 g of PbSO4. Assuming that all of the lead was precipitated, what is the percentage of Sn in the sample? (molar mass of PbSO4 = 303.3 g/mol)
A) 33.3% Sn
B) 17.7% Sn
C) 50.0% Sn
D) 66.7% Sn
E) 2.00% Sn

96. A 1.000-g sample of a metal chloride, MCl2, is dissolved in water and treated with excess aqueous silver nitrate. The silver chloride that formed weighed 1.286 g. Calculate the atomic mass of M.
A) 222.8 g
B) 76.00 g
C) 152.0 g
D) 304.0 g
E) none of these

97. A 0.307-g sample of an unknown triprotic acid is titrated to the third equivalence point using 35.2 mL of 0.106 M NaOH. Calculate the molar mass of the acid.
A) 247 g/mol
B) 171 g/mol
C) 165 g/mol
D) 151 g/mol
E) 82.7 g/mol

98. The following reactions
ZnBr2(aq) + 2AgNO3(aq) ® Zn(NO3)2(aq) + 2AgBr(s)
KBr(aq) + AgNO3(aq) ® AgBr(s) + KNO3(aq)
are examples of
A) oxidation-reduction reactions.
B) acid-base reactions.
C) precipitation reactions.
D) oxidation-reduction and precipitation reactions.
E) none of these

99. In the reaction C(s) + O2(g) ® CO2(g) carbon is __________.
A) the reducing agent
B) the electron acceptor
C) reduced
D) the oxidizing agent
E) more than one of these

100. Diabetics often need injections of insulin to help maintain the proper blood glucose levels in their bodies. How many moles of insulin are needed to make up 45 mL of 0.0052 M insulin solution?
A) 4.6 ´ 10–4 mol
B) 5.0 ´ 10–3 mol
C) 1.7 ´ 10–4 mol
D) 6.0 ´ 102 mol
E) 2.3 ´ 10–4 mol

101. Which of the following statements is(are) true? Oxidation and reduction
A) cannot occur independently of each other.
B) accompany all chemical changes.
C) describe the loss and gain of electron(s), respectively.
D) result in a change in the oxidation states of the species involved.
E) a, c, and d are true

102. In the reaction Zn + H2SO4 ® ZnSO4 + H2, which, if any, element is oxidized?
A) zinc
B) hydrogen
C) sulfur
D) oxygen
E) none of these

103. Given the following reaction in acidic media:
Fe2+ + Cr2O72– ® Fe3+ + Cr3+
answer the following question: The coefficient for water in the balanced reaction is
A) 1.
B) 3.
C) 5.
D) 7.
E) none of these

104. A glass column is filled with mercury and inverted in a pool of mercury. The mercury column stabilizes at a height of 735 mm above the pool of mercury. What is the pressure of the atmosphere?
A) 0.697 atm
B) 0.735 atm
C) 0.967 atm
D) 1.03 atm
E) 194 atm

105. You and a friend have gas samples in open manometers as shown:

You have Hg(l) in your manometer and your friend has water. The height h is the same in both manometers. Which of the following statements is true?
A) Your sample of gas has the higher pressure.
B) Your friend's sample of gas has the higher pressure.
C) Both samples of gas have the same pressure.
D) There is not enough information to answer the question.
E) None of these is correct.

106. A gas sample is held at constant pressure. The gas occupies 3.62 L of volume when the temperature is 21.6°C. Determine the temperature at which the volume of the gas is 3.45 L.
A) 309 K
B) 281 K
C) 20.6 K
D) 294 K
E) 326 K

107. Gaseous chlorine is held in two separate containers at identical temperature and pressure. The volume of container 1 is 1.30 L and it contains 6.70 mol of the gas. The volume of container 2 is 2.20 L. How many moles of the gas are in container 2.
A) 11.3 mol
B) 19.2 mol
C) 0.427 mol
D) 3.96 mol
E) none of these

108. A gas sample is heated from –20.0°C to 57.0°C and the volume is increased from 2.00 L to 4.50 L. If the initial pressure is 0.125 atm, what is the final pressure?
A) 0.189 atm
B) 0.555 atm
C) 0.0605 atm
D) 0.247 atm
E) none of these

109. The valve between a 5-L tank containing a gas at 9 atm and a 10-L tank containing a gas at 6 atm is opened. Calculate the final pressure in the tanks.
A) 3 atm
B) 4 atm
C) 7 atm
D) 15 atm
E) none of these

110. You fill a balloon with 2.50 moles of gas at 28°C at a pressure of 1.20 atm. What is the volume of the balloon?
A) 4.79 L
B) 22.4 L
C) 51.5 L
D) 56.0 L
E) 61.8 L

111. Body temperature is about 308 K. On a cold day, what volume of air at 273 K must a person with a lung capacity of 2.00 L breathe in to fill the lungs?
A) 2.26 L
B) 1.77 L
C) 1.13 L
D) 3.54 L
E) none of these

112. An automobile tire is filled with air at a pressure of 30 lb/in2 at 25°C. A cold front moves through and the temperature drops to 5°C. Assuming no change in volume, what is the new tire pressure?
A) 6.0 lb/in2
B) 28 lb/in2
C) 32 lb/in2
D) 20. lb/in2
E) 4.0 lb/in2

113. Given reaction N2 + 3H2 ® 2NH3, you mix 1 mol each of nitrogen and hydrogen gases under the same conditions in a container fixed with a piston. Calculate the ratio of volumes of the container (Vfinal/Vinitial).
A) 0.67
B) 1.00
C) 1.33
D) 1.50
E) none of these

114. Which of the following is the best qualitative graph of P versus molar mass of a 1-g sample of different gases at constant volume and temperature?

115. Given a cylinder of fixed volume filled with 1 mol of argon gas, which of the following is correct? (Assume all gases obey the ideal gas law.)
A) If the temperature of the cylinder is changed from 25°C to 50°C, the pressure inside the cylinder will double.
B) If a second mole of argon is added to the cylinder, the ratio T/P would remain constant.
C) A cylinder of identical volume filled with the same pressure of helium must contain more atoms of gas because He has a smaller atomic radius than argon.
D) Two of the above.
E) None of the above.

Use the following to answer question 116:

Four identical 1.0-L flasks contain the gases He, Cl2, CH4, and NH3, each at 0°C and 1 atm pressure.

116. For which gas do the molecules have the smallest average kinetic energy?
A) He
B) Cl2
C) CH4
D) NH3
E) all gases the same

Use the following to answer questions 117-118:

A plastic bag is weighed and then filled successively with two gases, X and Y. The following data are gathered:
Temperature: 0.0°C (273 K)
Pressure: 1.00 atmosphere
Mass of empty bag: 20.77 g
Mass of bag filled with gas X: 24.97 g
Mass of 1.12 liters of air at conditions given: 1.30 g
Volume of bag: 1.12 liter
Molar volume at STP: 22.4 liters

117. The mass of 1.12 liters of gas Y is found to be 6.23 g. The density of gas Y is
A) 10.6 g/L
B) 5.56 g/L
C) 15.6 g/L
D) 0.200 g/L
E) 0.180 g/L

118. The bag is emptied and refilled, successively, with gases X and Y, this time at 1 atm pressure and a temperature 30°C higher. Assume that the volume of the bag is the same as before. Which one of the following statements is wrong?
A) The full bag contains fewer molecules of each gas than it did at 0.0°C.
B) The ratio of the density of gas Y to the density of gas X is the same as at 0.0°C.
C) The molar masses of the two gases are the same as they were at 0.0°C.
D) The mass of each gas filling the bag is now 303/273 times the mass held at 0.0°C.
E) The average velocity of the molecules of gas X at 30°C is higher than it was at 0.0°C.

119. When 0.72 g of a liquid is vaporized at 110°C and 0.967 atm, the gas occupies a volume of 0.559 L. The empirical formula of the gas is CH2. What is the molecular formula of the gas?
A) CH2
B) C2H4
C) C3H6
D) C4H8
E) none of these

120. Gaseous C2H4 reacts with O2 according to the following equation:
C2H4(g) + 3O2(g) ® 2CO2(g) + 2H2O(g)
What volume of oxygen at STP is needed to react with 1.50 mol of C2H4?
A) 4.50 L
B) 33.6 L
C) 101 L
D) 67.2 L
E) Not enough information is given to solve the problem.

121. A 3.31-g sample of lead nitrate, Pb(NO3)2, molar mass = 331 g/mol, is heated in an evacuated cylinder with a volume of 1.62 L. The salt decomposes when heated, according to the equation
2Pb(NO3)2(s) ® 2PbO(s) + 4NO2(g) + O2(g)
Assuming complete decomposition, what is the pressure in the cylinder after decomposition and cooling to a temperature of 300 K? Assume the PbO(s) takes up negligible volume.
A) 0.380 atm
B) 0.228 atm
C) 0.0342 atm
D) 1.38 atm
E) none of these

122. The purity of a sample containing zinc and weighing 0.198 g is determined by measuring the amount of hydrogen formed when the sample reacts with an excess of hydrochloric acid. The determination shows the sample to be 84.0% zinc. What amount of hydrogen (measured at STP) was obtained?
A) 0.152 L
B) 0.0330 g
C) 3.42 ´ 10–3 mole
D) 1.53 ´ 1021 molecules
E) 1.53 ´ 1021 atoms

123. What volume of H2O(g) measured at STP is produced by the combustion of 4.00 g of natural gas (CH4) according to the following equation?
CH4(g) + 2O2(g) ® CO2(g) + 2H2O(g)
A) 5.60 L
B) 11.2 L
C) 22.4 L
D) 33.6 L
E) 44.8 L

124. Calcium hydride combines with water according to the equation
CaH2(s) + 2H2O(l) ® 2H2(g) + Ca(OH)2(s)
Beginning with 84.0 g of CaH2 and 36.0 g of H2O, what volume of H2 will be produced at 273 K and a pressure of 1520 torr?
A) 22.4 L
B) 44.8 L
C) 89.6 L
D) 179 L
E) none of these

125. A 1.00-g sample of a gaseous compound of boron and hydrogen occupies 0.820 L at 1.00 atm and 3°C. What is the molecular formula for the compound?
A) BH3
B) B2H6
C) B4H10
D) B3H12
E) B5H14

126. Given the equation
2KClO3(s) ® 2KCl(s) + 3O2(g)
A 3.00 g sample of KClO3 is decomposed and the oxygen at 24.0°C and 0.982 atm is collected. What volume of oxygen gas will be collected assuming 100% yield?
A) 304 mL
B) 608 mL
C) 911 mL
D) 1820 mL
E) none of these

Use the following to answer question 127:

Zinc metal is added to hydrochloric acid to generate hydrogen gas and is collected over a liquid whose vapor pressure is the same as pure water at 20.0°C (18 torr). The volume of the mixture is 1.7 L and its total pressure is 0.810 atm.

127. What would happen to the average kinetic energy of the molecules of a gas sample if the temperature of the sample increased from 20°C to 40°C?
A) It would double.
B) It would increase.
C) It would decrease.
D) It would become half its value.
E) Two of these.

128. A vessel with a volume of 10.0 L contains 2.80 g of nitrogen gas, 0.403 g of hydrogen gas, and 79.9 g of argon gas. At 25°C, what is the pressure in the vessel?
A) 0.471 atm
B) 6.43 atm
C) 3.20 atm
D) 5.62 atm
E) 2.38 atm

129. A balloon contains an anesthetic mixture of cyclopropane (cp) and oxygen (O2) at 170 torr and 570 torr, respectively. What is the ratio of the number of moles of cyclopropane to moles of oxygen?
ncp/no2 = ?
A) 0.19
B) 0.23
C) 0.30
D) 0.39
E) 0.46

130. Which of the following is not a postulate of the kinetic molecular theory?
A) Gas particles have most of their mass concentrated in the nucleus of the atom.
B) The moving particles undergo perfectly elastic collisions with the walls of the container.
C) The forces of attraction and repulsion between the particles are insignificant.
D) The average kinetic energy of the particles is directly proportional to the absolute temperature.
E) All of these are postulates of the kinetic molecular theory.

131. Use the kinetic molecular theory of gases to predict what would happen to a closed sample of a gas whose temperature increased while its volume decreased.
A) Its pressure would decrease.
B) Its pressure would increase.
C) Its pressure would hold constant.
D) The number of moles of the gas would decrease.
E) The average kinetic energy of the molecules of the gas would decrease.

132. Calculate the ratio of the effusion rates of N2 and N2O.
A) 0.637
B) 1.57
C) 1.25
D) 0.798
E) 1.61

133. A gas absorbs 0.0 J of heat and then performs 15.2 J of work. The change in internal energy of the gas is
A) –24.8 J
B) 14.8 J
C) 55.2 J
D) –15.2 J
E) none of these

134. Calculate the work for the expansion of CO2 from 1.0 to 2.5 liters against a pressure of 1.0 atm at constant temperature.
A) 1.5 liter × atm
B) 2.5 liter × atm
C) 0
D) –1.5 liter × atm
E) –2.5 liter × atm

135. Of energy, work, enthalpy, and heat, how many are state functions?
A) 0
B) 1
C) 2
D) 3
E) 4

136. Which of the following statements correctly describes the signs of q and w for the following exothermic process at P = 1 atm and T = 370 K?
H2O(g) ® H2O(l)
A) q and w are negative.
B) q is positive, w is negative.
C) q is negative, w is positive.
D) q and w are both positive.
E) q and w are both zero.

137. For a particular process q = 20 kJ and w = 15 kJ. Which of the following statements is true?
A) Heat flows from the system to the surroundings.
B) The system does work on the surroundings.
C) DE = 35 kJ.
D) All of the above are true.
E) None of the above are true.

138. Which statement is true of a process in which one mole of a gas is expanded from state A to state B?
A) When the gas expands from state A to state B, the surroundings are doing work on the system.
B) The amount of work done in the process must be the same, regardless of the path.
C) It is not possible to have more than one path for a change of state.
D) The final volume of the gas will depend on the path taken.
E) The amount of heat released in the process will depend on the path taken.

139. Which one of the following statements is false?
A) The change in internal energy, DE, for a process is equal to the amount of heat absorbed at constant volume, qv.
B) The change in enthalpy, DH, for a process is equal to the amount of heat absorbed at constant pressure, qp.
C) A bomb calorimeter measures DH directly.
D) If qp for a process is negative, the process is exothermic.
E) The freezing of water is an example of an exothermic reaction.

140. Two metals of equal mass with different heat capacities are subjected to the same amount of heat. Which undergoes the smallest change in temperature?
A) The metal with the higher heat capacity.
B) The metal with the lower heat capacity.
C) Both undergo the same change in temperature.
D) You need to know the initial temperatures of the metals.
E) You need to know which metals you have.

141. You take 200. g of a solid at 30.0°C and let it melt in 400. g of water. The water temperature decreases from 85.1°C to 30.0°C. Calculate the heat of fusion of this solid.
A) 125 J/g
B) 285 J/g
C) 461 J/g
D) 518 J/g
E) cannot without the heat capacity of the solid

142. Consider a rigid insulated box containing 20.0 g of He(g) at 25.0°C and 1.00 atm in one compartment and 20.0 g of N2(g) at 115.0 °C and 2.00 atm in the other compartment. These compartments are connected by a partition which transmits heat. What will be the final temperature in the box at thermal equilibrium? (Cv(He)= 12.5 J/K mol, Cv(N2)= 20.7 J/K mol)
A) 42.2°C
B) 58.9°C
C) 70.0°C
D) 81.0°C
E) none of these

143. Which of the following properties is (are) intensive properties?
I. mass
II. temperature
III. volume
IV. concentration
V. energy
A) I, III, and V
B) II only
C) II and IV
D) III and IV
E) I and V

144. Consider the reaction
H2(g) + (1/2)O2(g) ® H2O(l) DH° = –286 kJ
Which of the following is true?
A) The reaction is exothermic.
B) The reaction is endothermic.
C) The enthalpy of the products is less than that of the reactants.
D) Heat is absorbed by the system.
E) The reaction is exothermic and the enthalpy of the products is less than that of the reactants.

145. What is the heat capacity of mercury if it requires 167 J to change the temperature of 15.0 g mercury from 25.0°C to 33.0°C?
A) 6.92 ´ 10–3 J/g°C
B) 1.12 ´ 10–2 J/g°C
C) 0.445 J/g°C
D) 1.39 J/g°C
E) 313 J/g°C

146. A 4.0-g sample of Colorado oil shale is burned in a bomb calorimeter, which causes the temperature of the calorimeter to increase by 5.0°C. The calorimeter contains 1.00 kg of water (CH2O = 4.184 J/g°C) and the heat capacity of the empty calorimeter is 0.10 kJ/°C. How much heat is released per gram of oil shale when it is burned?
A) 21 kJ/g
B) 42 kJ/g
C) 0 kJ/g
D) 5.4 kJ/g
E) 5.2 kJ/g

147. Consider the following processes:
2A ® 1/2B + C DH1 = 5 kJ/mol
(3/2)B + 4C ® 2A + C + 3D DH2 = –15 kJ/mol
E + 4A ® C DH3 = 10 kJ/mol
Calculate DH for: C ® E + 3D
A) 0 kJ/mol
B) 10 kJ/mol
C) –10 kJ/mol
D) –20 kJ/mol
E) 20 kJ/mol

148. Consider the following processes:
DH (kJ/mol)
(1/2)A ® B 150.
3B ® 2C + D –125.
E + A ® D 350.
Calculate DH for: B + D ® E + 2C
A) 325 kJ/mol
B) 525 kJ/mol
C) –175 kJ/mol
D) –325 kJ/mol
E) none of these

149. Consider the following numbered processes:
I. A ® 2B
II. B ® C + D
III. E ® 2D
DH for the process A ® 2C + E is
A) DH1 + DH2 + DH3
B) DH1 + DH2
C) DH1 + DH2 – DH3
D) DH1 + 2DH2 – DH3
E) DH1 + 2DH2 + DH3

150. Acetylene (C2H2) and butane (C4H10) are gaseous fuels. Determine the ratio of energy available from the combustion of a given volume of acetylene to butane at the same temperature and pressure using the following data:
The change in enthalpy of combustion for
C2H2(g) = –49.9 kJ/g.
The change in enthalpy of combustion for
C4H10 = –49.5 kJ/g.

Answer Key -- honors final review

1. C
2. A
3. A
4. A
5. C
6. D
7. A
8. E
9. D
10. E
11. B
12. D
13. A
14. B
15. D
16. E
17. C
18. D
19. B
20. C
21. E
22. D
23. A
24. B
25. A
26. C
27. B
28. C
29. B
30. A
31. D
32. A
33. B
34. A
35. C
36. A
37. C
38. B
39. D
40. D
41. B
42. B
43. A
44. B
45. D
46. B
47. D
48. B
49. C
50. A
51. C
52. B
53. D
54. A
55. D
56. B
57. D
58. B
59. B
60. A
61. C
62. False
63. B
64. C
65. E
66. B
67. A
68. A
69. C6H12O2
70. C
71. D
72. D
73. D
74. D
75. C
76. B
77. B
78. E
79. B
80. B
81. B
82. A
83. A
84. C
85. B
86. A
87. A
88. D
89. C
90. B
91. C
92. E
93. B
94. A
95. A
96. C
97. A
98. C
99. A
100. E
101. E
102. A
103. D
104. C
105. A
106. B
107. A
108. E
109. C
110. C
111. B
112. B
113. A
114. A
115. E
116. E
117. B
118. D
119. C
120. C
121. A
122. D
123. B
124. A
125. B
126. C
127. B
128. D
129. C
130. A
131. B
132. C
133. D
134. D
135. C
136. C
137. C
138. E
139. C
140. A
141. C
142. A
143. C
144. E
145. D
146. D
147. C
148. C
149. D
150. About 2.21 times the volume of acetylene is needed to furnish the same energy as a given volume of butane.