Download Chemistry I - Net Start Class

Honors Chemistry
Fall Semester Exam Review
1. Which of the following is a physical property of magnesium?
is bendable
oxidizes to produce a white powder
melts at 650oC
silver colored metal
2. Which of the following is a physical change?
Water changes to steam.
A mixture of baking soda and vinegar give off a gas.
Heated moth balls turn into a liquid.
A powder mixed with water makes lemonade.
3. A solid, silver-colored metallic element is combined with a light green, gaseous element,
producing a powdery white solid product. What type of substance was the reactant?
4. A solid, pure, green substance was heated and yielded a white gas and a purple solid. The
material that was heated was what type of substance?
5. A copper-colored wire changes to a darker color after it is heated and then cooled. What type
of change has likely taken place to cause this color change?
6. In a proposal for a $1 million grant, a claim is made that a method will be developed to make
20 grams of gold from 10 grams of gold and no other ingredients. Is this possible?
7. In a patent application for a machine, 20 joules of energy is used to produce 35 joules of
energy. Is this possible?
8. At constant temperature, the volume of a fixed mass of a gas is inversely proportional to its
pressure. This statement is a(n) _______.
9. Can salt (NaCl) be chemically decomposed?
10. Acids behave in a particular manner because they donate a proton to a base. This statement
is a(n) _________.
11. Consider the chemical reaction in which carbon reacts with oxygen to produce carbon
dioxide. What mass of carbon dioxide would be produced if 24 grams of carbon reacted
completely with 64 grams of oxygen?
12. When paper becomes yellow-brown in color upon exposure to sunlight, what type of change
is likely taking place?
13. What state of matter is characterized by having an indefinite shape and volume and is
composed of charged particles?
14. Which state of matter takes both the shape and volume of its container and is composed of
neutral particles?
15. Which state of matter takes the shape and size of its container?
16. During condensation, heat is ______.
17. During vaporization, heat is ____.
18. During melting, heat is _____.
19. Which of the following is an example of matter?
air
heat
smoke
water
water vapor
20. What happens to the individual atoms in a chemical reaction?
21. How will the mass of a rusted iron nail compare with the mass of the same nail before it
rusted?
22. Under a certain set of conditions, the density of water is 1.00 g/ml and that of oxygen is 1.20
g/l. Which will be at the lowest level?
23. Calculate the average of 4.69 cm + 3.72 cm + 4.29 cm.
24. The temperature, in oC, of absolute zero is _____.
25. The sum of 0.223 m + 41.0 cm + 32.5 mm is ____.
26. 175 mg = ___ kg
27. 0.00150 km = ____ mm
28. What is the volume of each of the following:
a. 500 g of carbon tetrachloride D = 1.6 g/ml
b. 500 g of benzene D = 0.88 g/ml
c. 500 g of n-heptane D =0.68 g/ml
29. Three different people weigh a standard mass of 2.00 g on the same balance. Each person
obtains a reading of 7.32 g for the mass of the standard. These results imply that the balance that
was used is
a. accurate but not precise
b. precise but not accurate
c. accurate and precise
d. neither accurate nor precise
30. If the temperature of a piece of steel decreases, what happens to its density?
31. A student estimated a mass to be 250.g but, upon carefully measuring it, found the value to
the 240.g. What is the percent error of the estimated mass if the measured value is the accepted
one?
32. A beaker's mass is 21.05 g, while the mass of this beaker containing a sample of water is
22.15 g. What is the mass of the water sample?
33. An atom with a -1 charge has atomic number 19 and mass number 39. It has ___ electrons.
34. An atom has a charge of +1 and has 10 neutrons with a mass number of 21. What is its
atomic number?
35. A neutral atom has a mass number of 24 and 11 protons. How many neutrons does it have?
36. A neutral atom of bromine has a mass number of 80. What is its nuclear composition?
37. The average atomic mass of hydrogen is 1.008 amu. Which isotope of hydrogen
predominates - H-1, H-2, or H-3?
38. Indicate which pair(s) which are not isotopes;
atomic number
u
1
w
2
x
6
y
1
z
7
atomic mass
2
4
14
3
14
39. The element with atomic number 117 will have properties like those of
P
S
Cl
Ar
Na
40. Which of the following elements have properties closest to those of arsenic (As)?
Ge
P
Se
Te
41. The element with atomic number 116 will have properties like those of
P
S
Cl
Ar
Na
42. The element with atomic number 118 will have properties like those of
P
S
Cl
Ar
Na
43. What is the order of filling for the following orbitals?
4f
6s
5p
6p
5d
44. How many dots would appear in the Lewis electron dot diagram for an atom whose electron
configuration ended in 6s2 5d2 4f14 ?
45. What is the abbreviated electron configuration for the Zn2+ ion?
46. What is the electron configuration for silver?
47. What is the electron configuration of copper?
48. The first ionization energy of calcium is 589.9 kilojoules/mole. Which of the following
elements has a first ionization energy lower than that of calcium?
K
Sr
Rb
Mg
49. In which of the following pairs of particles does the first particle listed have the larger
radius?
Rb, Y
S -2, S
Br, F
Ga +3, As -3
50. Which of the following lists the elements in order of decreasing electronegativity?
O, P, Ge
Li, Na, K
Br, Cl, F
Cl, S, P
51. Phosphorus has a first ionization energy of 1011.8 kilojoules per mole. Which of the
following elements has a first ionization energy less than that of phosphorus?
F
N
Ga
Cl
Select the correct formula for:
52. ammonium sulfate
53. tin (IV) sulfide
54. ferric perbromate
55. periodic acid
56. calcium phosphide
57. phosphorous acid
58. cuprous carbonate
59. potassium hypobromite
60. tin (IV) sulfide
61. cuprous oxide
62. copper (II) sulfide
63. stannic oxide
64. True or False - Sodium monochloride is a valid name for NaCl.
65. True or False - Calcium (II) carbonate is a valid name for CaCO3.
66. Suppose there is an polyatomic ion, XO4-2, with the name clapperate. What would be the
formula of the polyatomic ion perclapperate?
67. Suppose there is an polyatomic ion, XO4-2, with the name clapperate. What would be the
formula of the polyatomic ion hypoclapperite?
68. True or False - Salt (NaCl) from the ocean and salt (NaCl) from a neutralization reaction are
the same.
69. True or False - Vitamin C that is produced in a factory and vitamin C from oranges is the
same.
70. When an ionic bond is formed, electrons are transferred from ______ to ______.
71. What type of compound will conduct an electric current when it is dissolved in water?
72. Draw the Lewis structure of the following compounds:
Cl2
CO2
N2
CO
F2
O2
H2
73. Which of the following statements was not part of Dalton’s hypothesis on the structure of
matter?
a. All matter is made of atoms.
b. Atoms of the same element are identical.
c. Atoms are made of protons and electrons.
d. Atoms unite in definite ratios to form compounds.
74. Which of these statements is NOT true?
a. atoms of the same element can have different masses
b. atoms of isotopes of an element have different numbers of protons
c. the nucleus of an atom has a positive charge
d. atoms are mostly empty space
75. (NH4)2SO4 has what type(s) of bonds?
76. Which of the these is a characteristic of most covalent compounds?
they are solids
when melted they conduct an electric current
they have low melting points
they are composed of metallic and nonmetallic elements
77. Which of the following has both covalent and ionic bonds?
CO
KC2H3O2
HCl
CoF2
KCl
78. What is the basis of a metallic bond?
79. Which of the following crystals is the most malleable?
ammonium nitrate
copper
aluminum phosphate
sodium iodide
80. What type of covalent bond is formed between a silicon atom and an oxygen atom?
81. What characteristic of metals makes them good electrical conductors?
82. Which of the following compounds is the most polar?
NH3
CBr4
I2
CO2
83. Which of the following covalent bonds is the least polar?
C-Br
C-H
C-S
C-Cl
C-C
84. What causes dipole interactions?
85. Under what conditions can potassium bromide conduct electricity?
What is the geometry of each of the following molecules?
86. AsCl3
87. HCN
88. BCl3
89. AsCl3
90. SiO2
Indicate if the following are polar (A) or nonpolar (B):
91. HCN
92. BCl3
93. AsCl3
94. SiO2
Indicate the type of Van der Waals force that predominates:
a. London (dispersion)
b. dipole interaction
c. hydrogen
95. H2
96. NO
97. CHCl3
98. Which of the following samples contains the largest number of atoms?
1.0 g of H2
6.02 x 1023 atoms of Ne
1.0 mole of H2
0.50 mole of H2
99. Calculate the mass in grams/atom of one iron atom.
100. Determine the mass in grams of one molecule of fluorine.
101. The number of oxygen atoms in 0.75 moles of Al2(SO4)3 is ______.
102. A compound with an empirical formula of CH4 has a molecular mass of 128. What is the
molecular formula?
103. What is the molecular formula of diethyl oxalate, a solvent used in some perfumes, if the
empirical formula is C3H5O2 and the molecular mass is 146?
104. A compound has a simplest formula of CH2N and a molar mass of 84 g/mole. What is its
molecular formula?
105. Thompson measured the charge to mass ratio of the proton as 9.5825 x 104 C/g and the
charge of the proton as 1.6022 x 10-19 C. Using this information, what is the calculated mass of
the proton?
106. Thomson measure the charge to mass ratio of the electron as -1.7588 x 108 C/g. Millikan
measured the actual charge as -1.6022 x 10-19 C. Using this information, what is the calculated
mass of the electron?
107. Draw the hydrogen bonding between HF and NH3.
108. Draw the hydrogen bonding between H2O and NH3.
109. Draw the hydrogen bonding between CH3OH and CH3NH2.
110. Draw the hydrogen bonding between HF and CH3NH2
111. Explain why the normal boiling point of CCl4 is 77oC and that of CBr4 is 190oC.
112. Explain why the bond length between the two carbon atoms is shorter in C2H4 than C2H6.
113. Explain why the long lengths in SO3 are all identical and are shorter than a sulfur-oxygen
bond.
114. Calculate the mass of 8.50 x 1023 oxygen molecules.
115. A crude sample of Ag2O has a mass of 50.0 g. It contains 25.0 g of Ag2O. What is the
percent of silver in the crude sample?
116. If 1.54 grams of an oxide of osmium (Os) are analyzed and found to contain 1.15 grams of
osmium, what is the empirical formula of this compound?
117. Determine the empirical formula of a compound if a sample of the compound contains
19.45 g of manganese and 7.55 g of oxygen.
For questions 118 and 119, indicate which is an analogy for:
a. decomposition
c. displacement
b. double displacement
d. combination
118. Two dancing partners exchange partners - two new couples are formed
119. Two people are legally married to form a couple
Identify the types of equations:
120. C + H2O  CO + H2
121. CaO + H2O  Ca(OH)2
122. NaHCO3  Na2CO3 + H2O + CO2
123. HCl + Na2CO3  NaCl + H2O + CO2
Balance the following equations:
124. C2H6 + O2 ----->CO2 + H2O
125. C6H8O6 + O2 -----> CO2 + H2O
126. C2H5OH + O2 ----> CO2 + H2O
127. SiF4 + H2O —> H2SiF6 + H2SiO3
128. Cl2 + KOH  KCl + KClO3 + H2O
129 Al(ClO3)3 -----> AlCl3 + O2
130. PCl5 + H2O -----> H3PO4 + HCl
131. H2O + P4O10 -----> H3PO4
132. CaCO3 + H3PO4 -----> Ca3(PO4)2 + CO2 + H2O
Complete and balance the following equations:
133. Cl2 + NaI 
134. Al + HC2H3O2
135. H2O2 
136. Na2SO4 + Mg3(PO4)2 
137. H2O + Na 
138. KNO3 
139. CaCO3 + H3PO4 
Problems:
1. Calculate the average of 4.69 cm + 3.72 cm + 4.29 cm.
2. The sum of 0.223 m + 41.0 cm + 32.5 mm is ____.
3. 175 mg = ___ kg (use dimensional analysis)
4. 0.00150 km = ____ mm (use dimensional analysis)
5. Calculate the volume of 550 g of carbon tetrachloride. D = 1.6 g/ml