ICP Semester Exam Review – Chemistry Topics
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. Check the lecture notes and old homework assignments.
I would also recommend going through all the tests.
FORMAT:
Questions will include multiple-choice. You will need a calculator and 2 pencils for the Scantron form.
A periodic table will be provided.
The Nature of Science – Ch. 1
After reading cooking instructions that said to add salt to water before boiling it, Jose guessed that adding salt must
make the water boil at a higher temperature. He decided to test his idea by performing the following experiment.
Jose measured out 1 quart of distilled water and added to it 2 tablespoons of salt. He then brought the water to a boil
and measured its maximum temperature. Jose ran two more trials using 2 tablespoons of salt. He then ran 3 trials
each with 4 tablespoons of salt and 6 tablespoon of salt. For each trial in his experiment, Jose used 1 quart of
distilled water and the same pot and stove. The average temperature for 2 tablespoons of salt was 102.7C. The
average temperature for 4 tablespoons of salt was 105.4C. The average temperature for 6 tablespoons of salt was
107.1C. Since Jose knew that the boiling point of water is 100C, he concluded that adding salt to water does cause
it to boil at a higher temperature.
1. What was Joe’s hypothesis? 4. List the dependent and independent variables.
2. What was Joe’s conclusion? VOCAB: hypothesis, theory, scientific law, control,
3. What component of experimental design is missing? constants, independent variable, dependent variable
Measurement – Ch. 2
5. A sample of cork has a mass of 54 g and a volume of 10. 230 g = kg
3
225 cm . What is its density? 11. 4.35 m = cm
6. Osmium is the densest element with a density of 22.57 12. 89.6 mm = m
3
g/cm . Find the mass of a sample of osmium that 13. How many centimeters long is the proverbial “10-foot
3
occupies a volume of 6.5 cm . pole?”
7. A typical ant is 3 ___ (unit?) long. 14. How many milliliters are in a 2.0 quart jug of milk?
8. A typical man weighs 85 ___ (unit?). 15. Calculate density from the slope of a graph (see test).
9. A soda can contains 355 ___ (unit?) of soda. VOCAB: number, quantity
Matter – Ch. 8 & 9
Classify as solid, liquid, gas, or plasma (20-23).). 25. When atmospheric pressure increases, boiling point __.
16. A balloon inflated in a cooled store pops in a hot car. 26. Describe energy and particles during a phase change.
17. An air bubble released 30 feet underwater, expands in 27. Describe energy and particles between phase changes.
size as it rises to the surface. Classify the following as element, compound, solution, or
18. As temperature increases, the liquid in a thermometer heterogeneous mixture (36-39).
rises up the tube. 28. steam 30. liquid nitrogen
Identify which phase change is being described (27-32). 29. smoke 31. Kool-Aid®
19. A transition from gas to liquid. 32. Compare & contrast solutions, colloids, & suspensions.
20. A transition from liquid to gas at the boiling point. Classify the following as chemical or physical (41-48).
21. A transition from liquid to solid. 33. conducts electricity 37. decomposing plant
22. A transition from solid to gas. 34. explosive 38. sublimation of dry ice
23. A transition from liquid to gas below the boiling point. 35. corrosive 39. grating cheese
24. A transition from solid to liquid. 36. dissolves in water 40. acid rain damage to
marble
41. Compare and contrast mixtures and compounds.
VOCAB: kinetic molecular theory, thermal expansion,
kinetic energy, potential energy, absolute zero
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Atomic Structure & The Periodic Table – Ch. 10
Which scientist is responsible for each discovery (50-56)? 46. The largest atoms are in the ___ corner of the table. :
42. Developed the “electron cloud” model of the atom. Ba or Rn?
43. Proposed that electrons travel in circular orbits. 47. Draw the electron dot diagrams for Rb and S.
44. Draw the Bohr model diagram for Sodium VOCAB:, isotope, average atomic mass, valence electrons,
45. List the subatomic particles & isotope symbol for period, group, metals/nonmetals/metalloids, /transition
bromine-80. metals/inner trans metals
Chemical Bonds – Ch. 11
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48. Why do most atoms form bonds to get 8 valence e ? 57. Involve the unequal sharing of electrons.
Are these compounds ionic or covalent (49-51)? 58. Create partial + and – charges within the molecule.
49. SO3 KCl PbSO3 59. Involve the equal sharing of electrons.
50. FeCl3 H2CO3 MgBr2 60. Name these molecular compounds: PCl3, SO2, N2O5.
51. NaNO3 61. Name these ionic compounds: LiBr, MgSO4, NaNO2.
Identify these properties as ionic or covalent (52-55) 62. Name these acids: H2SO4, HClO2, HF.
52. Form individual molecules. 63. Write molecular formulas for: carbon tetrachloride,
53. Involve the transfer of electrons.. sulfur hexafluoride, dinitrogen monoxide.
54. Involve the sharing of electrons. 64. Write ionic formulas for: calcium chloride, aluminum
55. Conduct electricity in solution. oxide, copper(II) nitrate.
Identify these properties as polar or nonpolar (78-81). 65. Write acid formulas for: nitrous acid, phosphoric acid,
56. Usually formed between identical atoms. hydrobromic acid.
VOCAB: octet rule, polar, nonpolar, diatomic elements,
oxidation number, polyatomic ion
Solutions – Ch. 15
66. Describe the dissolving process (solvation).
67. What conditions cause solids to dissolve faster? 72. Addition of a solute to the solvent causes the boiling point to
68. What conditions cause gases to dissolve faster? ___ and the freezing point to ___.
69. If additional solute dissolves, the solution is ___. 73. Which solution will have a greater effect on freezing point:
70. If additional solute causes crystallization, the solution is ___. 20% NaCl or 30% NaCl?
71. If additional solute doesn’t dissolve, the solution is __. VOCAB: solute, solvent, solvation, solubility, concentrated/dilute,
unsaturated/saturated/supersaturated, detergent, electrolyte,
dissociation, ionization
Acids & Bases – Ch. 17
Identify these properties as acid, base, or both (13-20). 82. Identify common substances as acids or bases using pH
74. A corrosive electrolyte. value or litmus test.
75. Produces hydroxide ions in solution. 83. The reactants in a neutralization reaction are ___.
76. Sour taste. 84. The products of a neutralization reaction are ___.
77. Produces hydrogen ions in solution. 85. A salt is formed from the ___ of an acid and the ___ of a
78. Bitter taste and slippery feel. base.
79. Soap and ammonia are examples. VOCAB: acid, base, indicator, strong/weak acid or base, pH,
80. Can be detected with an indicator. buffer, neutralization reaction, salt
81. Vinegar and lemon juice are examples.
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IPC 1st Semester Exam Review – ANSWER KEY
1. applied heterogeneous mixtures that scatter light,
2. pure whereas solutions are homogeneous mixtures
3. applied that don’t scatter light. Solutions have the
4. pure smallest particles, suspensions have the largest
5. Adding salt makes the water boil at a higher particles.
temperature. 41. physical
6. Adding salt to water does cause it to boil at a 42. chemical
higher temperature. 43. chemical
7. Joe correctly tested a single variable and ran 44. physical
repeated trials, but he did not have a control 45. chemical
(water without salt) for comparison. 46. physical
8. dependent variable-boiling temperature, 47. physical
independent variable-amount of salt 48. chemical
9. 0.24 g/cm3 49. Both mixtures and compounds contain more
10. 146.7 g than one element. Mixtures can be physically
11. mm (millimeters) separated and have a random composition.
12. kg (kilograms) Compounds can only be chemically separated
13. mL (milliliters) and have a definite composition.
14. 0.23 kg 50. Dalton
15. 435 cm 51. Rutherford
16. 89,600 m 52. Becquerel
17. 304.8 cm 53. Thomson
18. 1892 mL 54. Chadwick
19. 0.8 g/cm3 55. Schrödinger
20. solid 56. Bohr
21. plasma 57. see Atomic Structure Timeline Lecture Handout
22. gas
23. liquid
24. Charles’ law (V&T)
12p
25. Boyle’s law (P&V) 12n
26. thermal expansion
27. condensation
28. vaporization (boiling) 58.
29. freezing 80
59. 35 Br , atomic# 35, mass# 80, 35 protons, 45
30. sublimation
31. evaporation neutrons, 35 electrons
32. melting 60. 6.92 u
33. increases 61. atomic mass, atomic number
34. Potential energy increases since temperature is 62. bottom-left, top-right
constant. The spacing between particles 63. Li
increases. 64. Xe
35. Kinetic energy increases since temperature is 65. Li
increasing. The particles move faster. 66. Rn
36. compound S
37. heterogeneous mixture 67. Rb
38. element 68. Having 8 valence electrons gives most atoms a
39. solution full outer energy level which makes the atoms
40. Solutions and colloids have particles that don’t more stable.
settle, whereas particles in a suspension do 69. covalent
settle. Colloids and suspensions are 70. ionic
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71. ionic
72. covalent
73. ionic
74. ionic
75. covalent
76. covalent
77. ionic
78. nonpolar
79. polar
80. polar
81. nonpolar
82. phosphorus trichloride, sulfur dioxide, dinitrogen pentoxide
83. lithium bromide, magnesium sulfate, sodium nitrite
84. sulfuric acid, chlorous acid, hydrofluoric acid
85. CCl4, SF6, N2O
86. CaCl2, Al2O3, Cu(NO3)2
87. HNO2, H3PO4, HBr
88. Two molecules of sulfur dioxide gas react with one molecule of
oxygen gas to form two molecules of sulfur trioxide gas.
89. requires, releases
90. endothermic
91. exothermic
92. endothermic
93. endothermic
94. exothermic
95. exothermic
96. decomposition: SiCl4 Si + 2Cl2
97. synthesis/combustion: 4Li + O2 2Li2O
98. combustion: C3H8 + 5O2 3CO2 + 4H2O
99. single replacement: 3H2SO4 + 2Al Al2(SO4)3 + 3H2
100. synthesis: 2Fe + 3Cl2 2FeCl3
101. double replacement: H2CO3+ Ba(NO3)2 2HNO3 + BaCO3
102. single replacement: Al2S3 + 3Cl2 2AlCl3 + 3S
103. double replacement: CuCl2 + 2AgNO3 Cu(NO3)2 + 2AgCl
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