# 2nd Semester Exam Review - Spring 2002 - DOC

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HONORS CHEMISTRY – FINAL EXAM REVIEW – SPRING
2010
STRATEGY: Start by reading through your notes to refresh your memory on these topics. Then, use this review sheet as a starting
point to identify the areas on which you need to spend more study time. For those areas, go back to homework assignments, quizzes,
and reviews to practice more problems. Look back at all your old tests from this semester!! Keep in mind that these questions are
only samples and do not include specific examples of how vocabulary and other conceptual information might appear in a scantron
format. Remember you can access notes on my website.

Liquids & Solids – Chapter 13
1. Compare and contrast liquids and solids.
2. Explain the relationship between strong intermolecular forces and the following properties –
volatility, vapor pressure, and boiling point.
Draw a heating curve and indicate whether it would be flat or rising in 3 – 7.
3. liquid is boiling                     6. potential energy is increasing
4. solid is warming                      7. kinetic energy is increasing
5. solid is melting
VOCAB:                     vaporization           crystalline vs. amorphous
evaporation            sublimation
volatility             endothermic
vapor pressure          exothermic
Gases – Chapter 14
Identify the gas law and solve the problem (8-21).
8. Hydrogen gas is collected over water at 35C to give a total pressure of 0.80 atm. Find the
pressure of the dry hydrogen gas in kPa. (see water vapor pressure table)
9. A jar is tightly sealed at 22C and 772 torr. What is the pressure inside the jar after it has been
heated to 178C?
10. 300.0 mL of gas has a pressure 75.0 kPa. When the volume is decreased to 125.0 mL, what is its
pressure?
11. 50.0 L of gas has a temperature of 75C. What is the temp in Celsius when the volume changes
to 110 L?
12. What is the volume of a container that holds 48.0 g of helium at a pressure of 4.0 atm and
temperature of 52C?
13. A gas occupies 325 L at 25C and 98.0 kPa. What is its volume at 70.0 kPa and 15C?
14. What volume of SO2 is produced from 32.5 g of ZnS at 23C and 103.3 kPa? (BALANCE!)
ZnS(s) + O2(g)           ZnO(s) + SO2(g)
15. A student adds 5.85 g of dry ice (solid CO2) to an empty balloon. What will be the volume of the
balloon at STP after all the dry ice sublimes?
16. What is the density of ammonia gas at STP?
17. 20.0 g each of helium and an unknown diatomic gas are combined in a 1500. mL container. If the
temperature is 298 K, and the pressure inside is 86.11 atm, what is the unknown gas?
18. A compound has the empirical formula C2H3. A 1000.-mL flask, at 31oC and 742 torr, contains
2.12 g of the gaseous compound. Give the molecular formula.
19. The partial pressure of N2(g) is 0.275 atm and that of O2(g) is 0.150 atm in a mixture of the two
gases.
a. What is the mole fraction of each gas in the mixture?
b. If the mixture occupies a volume of 7.36 L at 75oC, calculate the number of moles of gas in the
mixture.
c. Calculate the number of grams of each gas in the mixture.

20. A 460.8-g sample of X2(g) has a volume 9.0 L at 10 atm and 102 °C. What is element X?
21. A balloon contains 0.1 moles of oxygen and 0.4 moles of nitrogen. If the balloon is at standard
temperature and pressure, what is the partial pressure of the nitrogen?
22. Define real gases. When do they behave like ideal gases?

VOCAB:Kelvin             diffusion
STP                  effusion
Thermochemistry – Chapter 17
23. The specific heat capacity for silver is 0.24 J/goC. Calculate the energy required to raise the
temperature of 50.0 g Ag from 263 K to 300. K. Calculate the molar heat capacity of silver.
24. It takes 490. J of energy to raise the temperature of 125.6 g Pb from 20.0oC to 50.5oC. Calculate the
specific heat capacity and the molar heat capacity of Pb.
25. A 45.2-g sample of copper is heated to 96.4oC and then placed in a calorimeter containing 75.0 g water
at 18.6oC. The equilibrium temperature in the calorimeter is 20.8oC. Calculate the specific heat
capacity of copper; assuming that all the heat lost by the copper is gained by the water.
26. A 25.0-g sample of nickel metal is heated to 100.0oC and dropped into 45.0 g of water, initially at
22.5oC. Assuming that al the heat lost by the nickel is absorbed by the water; calculate the final
temperature of the nickel and the water. The specific heat of nickel is 0.444 J/goC.
27. For the combustion of propane, C3H8 calculate the enthalpy change for the burning of 10.00 g of
propane.
28. Use Hess’s Law to determine H for this reaction: 2Al + Fe2O3  2Fe + Al2O3
Given these two reactions:         4Al + 3O2  2Al2O3         H = -1233.7 kJ
4Fe + 3O2  2Fe2O3         H = -1648.5 kJ

Structure of the Atom – Chapter 4
29. Identify the scientists who made the following discoveries.
a. Atoms contain negative particles called electrons.
b. Developed the atomic theory.
c. Atoms contain a dense, positive nucleus.
d. Atoms are indivisible.
30. Describe the evolution of the atomic model from the Democritus to the electron cloud model.
31. Write the isotope symbol, including atomic number & mass number, for the following isotopes.
a. carbon-14                           c.    nickel-63
b. chromium-53                         d.    zirconium-92
32. Complete the table for the following isotopes.
Symbol                Zn
Atomic #                        20
Mass #                65                   74         40
# of protons                               34
# of neutrons                   21
# of electrons                                        18
33. Calculate the average atomic mass of copper if 69.17% of the copper atoms occurring in nature
are 63Cu and 30.83% are 65Cu.

VOCAB: isotope
average atomic mass
2
quark
Electrons in Atoms – Chapter 5
34. Calculate the wavelength of a photon if its frequency is 2.5 x 105 Hz.
35. Find the energy of a photon if its frequency is 7.31 x 1014 Hz.
36. What is the primary difference between the modern model of the atom and Bohr’s model?
37. Draw orbital diagrams for the following elements.
Symbol Atomic # Orbital Diagram
F

V
38. Explain why chromium’s electron configuration is [Ar] 4s13d5 instead of the expected configuration
of [Ar] 4s23d4.
39. Give the noble gas electron configuration for the following.
Noble Gas e-
Symbol # e-
Configuration
Pd

At

VOCAB:
excited state/ground state                               photon
wave-particle duality                                    Hund’s Rule
valence/core e-                                          Aufbau Principle
Pauli Exculsion Principle                                Heisenberg Uncertainty Principle
Periodic Table and Periodic Law – Chapter 6
40. How did Mendeleev and Moseley arrange the elements in the periodic table?
41. Circle the atom with the LARGER radius.
a. Ra      N
b. Ne      Xe
42. Circle the atom with the LARGER radius.
a. Cl      S
b. Mg      Be
43. Circle the atom with the HIGHER first ionization energy.
a. Li      Cs
b. Ba      As
44. Circle the atom with the HIGHER electronegativity.
a. Cl      Si
b. Cs      K
45. Why is there a large increase in ionization energy when the 4th electron is removed from
aluminum?

VOCAB:ionization energy
periodic law
metals/nonmetals/metalloids
shielding
electronegativity
3
Ionic Compounds and Metallic Bonding – Chapter 7
46. Based on their electronegativities, are the bonds in the following substances IONIC, POLAR, or
NONPOLAR?
a. MgO                          c.     LiCl
b. H2O                          d.     Br2
47. Are the following properties characteristics of ionic or covalent bonding?
a. These bonds involve a transfer of electrons.
b. Substances containing these bonds do not conduct electricity and have low melting points.
c. Compounds containing these bonds have a crystal lattice structure.
d. These bonds are formed by sharing electrons.
48. Write formulas for the following compounds (HINT: First determine ionic/acid/covalent).
a. calcium chloride                           d.      sulfur trioxide
b. iron(III) nitrate                          e.      dinitrogen pentoxide
c. hydrobromic acid                           f.      sulfurous acid
49. Write names for the following compounds (HINT: First determine ionic/acid/covalent).
a. CrCl2                                      d.      CaSO4
b. Fe2CO3                                     e.      P5O8
c. As2Cl3                                     f.      HClO4
50. Explain the difference between nonpolar covalent, polar covalent, and ionic bonds in terms of
sharing of electrons and electric charge.

VOCAB:bond energy (bond length)
chemical bond
potential energy

Covalent Bonding – Chapter 8
51. For each of the compounds, complete all the information:

MOLECULE       LEWIS       e- TALLY     HYBRID-    ELECTRON     MOLECULAR       BOND       POLAR OR
DIAGRAM                   IZATION    GEOMETRY     GEOMETRY        ANGLE      NONPOLAR

a. SeO3

b. AsH3

c. SO4 2-

d. BeCl2

e. PCl3

f. SCl6

g. H2O

52. For each of the molecules above, determine all the intermolecular forces that are present.
53. What is a resonance structure?
54. Draw the Lewis structure for SO2.

4
Mixtures & Solutions – Chapters 15 & 16
55. Explain the effect of adding more solute to unsaturated, saturated, and supersaturated solutions.
56. Explain how temperature and polarity affect solubility.

State whether each pair is soluble or insoluble (57-60).
57. KCl in water          59. wax in C6H6
58. ammonia in oil        60. CH4 in water
61. How many grams of AlCl3 are required to make a 2.25m solution in 30.0 g of water?
62. Find the molarity of a 750 mL solution containing 346 g of potassium nitrate.
63. Calculate the number of grams required to make a 50.0 mL solution of 6.0M NaOH.
64. What volume of 12M HCl is needed to prepare 250 mL of 0.20M HCl?
65. How many grams of chromium (III) are required to react with 125 mL of 0.75M CuSO4?
66. Explain the difference in preparing solutions based on molarity versus molality.
67. Which will have the greatest effect on tf at the same molality: C12H22O11, MgBr2, AlCl3, or
NH4NO3? Why?
68. When 26.4 g of NaBr dissolves in 0.20 kg of water, what is the freezing point of the solution?
VOCAB:solvation              solubility    ionization    molality     strong/weak/nonelectrolyte
Acids and Bases – Ch. 19
State whether the following are acids or bases (69-72).
69. Have a sour taste                          71. Feel slippery
70. React with metals                          72. Turn blue litmus paper red
73. Define acids and bases according to Arrhenius, Brønsted-Lowry, and Lewis.
74. Identify each substance as acid, base, conjugate acid, or conjugate base.
H2S + H2O  HS – + H3O+
75. Give the conjugate acids of: NH3 and Br –.
76. Give the conjugate bases of: H3O+ and HSO4–.
77. Find the pH of 0.75M HCl. Is this acidic or basic?
78. Find the molarity of a KOH solution with a pH of 9.5.
79. Is the solution in #78 acidic or basic?
80. What is the pH of a solution of 16.67 mL of 2.000M hydrochloric acid in 1900. mL of water?
81. When a neutralization reaction between a strong acid and a weak base reaches the equivalence
point, will the solution be acidic, basic, or neutral? Sketch a possible titration curve.
82. If 43.5 mL of 0.15 M HBr is required to neutralize 25.0 mL of Ca(OH)2, what is the molarity of
Ca(OH)2?
83. What volume of 0.500M phosphoric acid is required to neutralize 50.0 mL of 0.150M magnesium
hydroxide?
VOCAB:                        hydronium ion                 neutralization reaction
amphoteric substance          titration
strong/weak acid/base         equivalence point
buret                         indicator
Reaction Rates & Equilibrium – Chapter 18
84. List the 5 factors that affect the rate of reaction. How do each of these factors specifically affect
the rate of reaction?
85. In order to have an “effective” collision, what 3 things must be present?
86. What is Ea?
87. Draw an energy diagram that depicts an exothermic reaction.
88. Label the products, reactants, Ea, and activated complex for the diagram you drew for #87.
89. What makes the diagram in #87 exothermic?
5
90. Is the reaction below endothermic or exothermic? How do you know?
2H2 + 2O2  2H2O + 571.6 kJ
91. What is the law of mass action?
92. What is a reaction quotient (Q)?
93. Write the equilibrium expression for the following equation:
HNO3(aq) + NH3(g) ↔ NH4NO3(aq)

94. Write the equilibrium expression for the following equation:
SiO2(s) + 4HF(g) ↔ SiF4(g) + 2H2O(g)

95. Write the equilibrium expression for the following equation:
CO(g) + H2(g) ↔ C(s) + H2O(g)

96. Write the equilibrium expression for the following equation:
CH4(g) + 2H2S(g) ↔ CS2(g) + 4H2(g)

97. Complete the chart below (right, left, decreases, increases, remains the same, etc.) for the
following equation:  2HCl(g) + 49.7 kJ ↔ H2(g) + Cl2(g)

Equilbrium
Stress                                [H2]             [Cl2]            [HCl]              K
Shift
Remove H2
Remove Cl2
Remove HCl
Increase Temp
Decrease Temp
Increase Pressure
Decrease Pressure

98. For the reaction: H2(g) + I2(g)      2HI(g), Keq = 51. Using the concentrations given, predict how
the reaction will proceed to reach equilibrium if we start with 2.0 × 10-2 mol of HI, 1.0 × 10-2 mol of
H2, and 3.0 × 10-2 mol of I2 in a 2.0-L container.

99. For the reaction below, the mixture of gases was found to have concentrations of 0.1207 M H2,
0.0402 M N2, and 0.00272 M NH3. Predict the direction that the reaction will proceed if the
equilibrium constant, Keq = 0.105.

100.   Explain the relationship between entropy, free energy and spontaneity in reactions.

VOCAB:          activated complex         catalyst              activation energy

6
HONORS CHEMISTRY – FINAL EXAM REVIEW – SPRING 2010 ANSWER KEY

1.    Both are incompressible with high density. Liquids are fluids. Solids have stronger IMF and slower diffusion.
2.    Strong IMF means molecules want to stay in the liquid state so volatility is low. Since there are fewer vapor molecules, v.p. is low.
The b.p. is high because higher temps are needed to overcome the strong forces.
3.    flat                          6. flat
4.    rising                        7. rising
5.    flat
8.    Dalton, 0.74 atm, 560 mm Hg, 76 kPa
9.    Gay-Lussac, 1180 torr
10.   Boyle, 180. kPa
11.   Charles, 490C
12.   Ideal, 80. L
13.   Combined, 440. L
3                                        3
14.   7.95 dm SO2 (or 7.95 L SO2) (Remember: dm = L)
15.   Ideal, 2.98 L
16.   Ideal, 0.76 g/L
17.   Ideal, MM = 70.9g/mol, gas = Cl2
18.   Ideal, MF = C4H6
19.   a. χN2 = 0.647, χO2 = 0.353; b. n = 0.110 moles of gas; c. 1.99 g N2, 1.24 g O2
20.   MM = 78.8 g/mol, X = Br
21.   PN2 = 0.8 atm
22.   Real gas molecules have a volume and attract each other. They act ideal at high temperatures and low pressures.
o
23.   q = 440 J, molar heat capacity of Ag = 25.9 J/mol C
o                 o
24.   c for Pb = 0.128 J/g C and 26.5 J/mol C
o
25.   c for Cu = 0.202 J/g C
o
26.   TF = 41.3 C
27.   -503.3 kJ
28.   H = 207.4 kJ
29.   a. Thomson, b. Dalton, c. Rutherford, d. Democritus
30.   Dalton’s billiard ball model-sphere of uniform density. Thomson’s plum pudding model-negative electrons dispersed in positive
atom. Rutherford’s nuclear model-dense, positive nucleus surrounded by negative electrons. Bohr’s planetary model-electrons
move in circular orbits in specific energy levels. Schrödinger’s electron cloud model-electrons move within orbitals not in specific
orbits. (Chadwick then added neutrons to the nucleus.)
14    53     63    92
31.    6 C, 24 Cr, 28Ni, 40 Zr
32.    Symbol           Zn        Ca         Se        Ar
Atomic #         30        20         34        18
Mass #           65        41         74        40
# of protons     30        20         34        18
# of neutrons    35        21         40        22
# of electrons 30          20         34        18
33.   63.62 amu
34.   1200 m
4.84  10 J
-19
35.
36.   Bohr’s model stated that electrons circled the nucleus in fixed, circular paths called orbits. The modern model states that electrons
move around the nucleus in orbitals where there is a probability of finding an electron.
37.
F    9                                          V    23

38. In order to achieve greater stability, Cr moves one electron from the 4s-sublevel to the 3d-sublevel to make it half-full.
2   8
39.    Pd 46         [Kr] 5s 4d
2 14   10  5
At     85     [Xe] 6s 4f 5d 6p
40.   Mendeleev arranged the elements in order of increasing               47. a. ionic, b. covalent, c. ionic, d. covalent
atomic mass. Mosely arranged them by increasing atomic               48. a. CaCl2, b. Fe(NO3)3, c. HBr, d. SO3, e. N2O5, f.
number.                                                                  H2SO3
41.   a. Ra, b. Xe                                                         49. a. chromium(II) chloride, b. iron(I) carbonate, c.
42.   a. S, b. Mg                                                              diarsenic trichloride, d. calcium sulfate, e.
43.   a. Li, b. As                                                             pentaphosphorous octoxide, f. perchloric acid.
-
44.   a. Cl, b. K                                                          50. nonpolar covalent – e are shared equally, symmetrical
th
45.   Removing the 4 electron from aluminum represents                         orbital overlap, no separation of charge
-
removing a core electron.                                                polar covalent – e are shared unequally, lopsided overlap,
46.   a. ionic, b. polar, c. ionic, d. nonpolar                                partial charges
7
-
ionic – e are not shared, no overlap, complete charges               87. See exothermic diagram labeled #87
51.                                                                      88. see labels
Molecule    Lewis     e-    Hybrid    E-        Molec Bond P/            89. Energy is released by the reaction
Geo       Geo       Angle N        90. Exothermic – energy is released (on the products side of
2
a. SeO3                24 sp          trig      trig      120       N        the reaction)
planar    planar                   91. States that…
b. AsH3                8    sp
3
tetrah    trig      109.5 P        92. comp of equil…
pyram                                     [ NH 4 NO3 ]
c. SO4
2-
32 sp
3
tetrah tetrah 109.5 N              93.   K eq 
d. BeCl2               16 sp          linear    linear    180       N                   [ HNO 3 ][ NH 3 ]
3
e. PCl3                26 sp          tetrah trig         109.5 P
[ H 2 O]2 [SiF4 ]
3 2
pyram                    94.   K eq 
f. SCl6                48 sp d
3
octah     octah     90        N                        [ HF ]4
g. H2O                 8    sp        tetrah bent         109.5 P
52. a. disp; b. disp, dipole, H; c. disp; d. disp; e. disp, dipole; f.                    [ H 2 O]2
disp; g. disp, dipole, H                                            95.   K eq   
53. one of two or more equally valid Lewis structures of a
[ HF ][ CO ]
molecule or polyatomic ion                                                          [ H 2 ]4 [CS 2 ]
54. diagram ->                                                           96.   K eq 
[CH 4 ][H 2 S ]2
97.
55. Unsaturated – solute will dissolve. Saturated – solute will          Stress         Eq         [H2]      [Cl2]    [HCl]     K
not dissolve. Supersaturated – rapid crystallization.                               Shift
56. Solubility of gases increases with low temps & high                  Add [H2]       L          -                          same
pressure. Solubility of solids increases with high temps.
Add [Cl2]      L                   -                 same
“Like dissolves like” describes solubility and polarity
relationship.                                                        Add HCl        R                           -         same
57. soluble (P/P)                  59. soluble (NP/NP)                   Remove         R          -                          same
58. insoluble (P/NP)               60. insoluble (NP/P)                  H2
61. 9.00 g AlCl3                                                         Remove         R                   -                 same
62. 4.6M KNO3                                                            Cl2
63. 12 g NaOH                                                            Remove         L                           -         same
64. 4.2 mL of 12M HCl                                                    HCl
65. 3.3 g Cr                                                             Increase       R                                    
66. Molarity – measure amount of solute, add enough water to             Temp
reach the desired volume. Molality – measure amount of               Decrease       L                                    
solute, measure kg of water, combine.                                Temp
67. C12H22O11 – 1, MgBr2 – 3, AlCl3 – 4, NH4NO3 – 2                      Increase       No         No        No       No        same
68. – 4.8C                                                              Press          effect     effect    effect   effect
69. acid                           71. base                              Decrease       No         No        No       No        same
70. acid                           72. acid                              Press          effect     effect    effect   effect
+                                –
73. Arr acid – forms H3O in water. Arr base – forms OH in
water. B-L acid – proton donor, B-L base – proton acceptor.          98. Q = 1.3, Q<Keq, shift toward products (right)
-                               -
Lewis acid – e pair acceptor, Lewis base – e pair donor.             99. Q = 0.105, at equilibrium
74. A, B, CB, CA                                                         100. increase in entropy and free energy released in a reaction
+
75. NH4 and HBr                                                              leads to spontaneous reactions
2–
76. H2O and SO4
77. 0.12, acidic
-5
78. 3.2 × 10 M KOH (pOH = 4.5)
79. basic
80. pH = 1.760
81. acidic, see curve labeled #81 below
82. 0.13M Ca(OH)2
83. 10.0 mL H3PO4
84. Surface area, temperature, concentration, nature of
reactants, catalyst
85. Collision, correct orientation, enough energy to form
activated complex                                                                                  #87
86. Activation energy – minimum amount of energy needed for
the reaction to occur

8

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