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Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.)
1.
Which unit is NOT used for volume?
a) L
b) mL
c) mm
d) cm3
2.
Which piece of equipment would be best to
measure 56.7 mL of water?
a) graduated cylinder
b) balance
c) test tube
d) electronic balance
3.
Which of these are the smallest?
a) 1 liter
b) 1microliter
c) 1milliliter
d) 1kiloliter
4.
The prefix kilo means:
a) 100 times larger
b) 1000 times larger
c) 1 000 000 times larger
d) 10 times smaller
5.
The metric prefix milli means:
a) 100 times larger
b) 1000 times larger
c) 1000 times smaller
d) 100 times smaller
6. The number 2 X 101 expressed in standard
notation is:
a) 200
b) 20
c) 2
d) 0.02
7. The number 3 X 10 -3 expressed in standard
notation is
a) 300
b) 0.300
c) 0.003
d) 3.000
8. How many of the zeros in the measurement
0.000040300 are significant?
a) 8
b) 6
c) 5
d) 3
9. The number of significant figures in the
measurement 0.070g is:
a) 1
b) 2
c) 3
d) 4
10. When adding and subtracting measurements one should
limit and round answers to:
a) the least number of significant figures in any of the
measurements
b) the least number of decimal places in any of the
measurements
c) the tenths place
d) three significant figures
11. When multiplying and dividing measurements, one
should limit and round answers to:
a) the least number of significant figures in any of the
factors
b) the least number of decimal places in any of the
measurements
c) the tenths place
d) three significant figures
12. When measuring the volume of a liquid in a graduated
cylinder, the measurement should be read from the:
a) top of the meniscus
b) top of the cylinder
c) bottom of the meniscus
d) table top
13. If the mass of a dry beaker is 19.02 grams and increases
to 22.40 grams when a sample is added, what is the
mass of the sample?
a) 22.40 g
b) 41.42 g
c) 3.38 g
d) 1.10 g
14. The mass of a sample is 4.11grams. The volume is
2.00cm3. What is its density?
a) 8.22g/cm3
b) 2.06g/cm3
c) 0.49g/cm3
d) not enough information
15. A student estimated a mass to be 250g, but upon
carefully measuring it, found it to be 240g, What is the
percent error of the estimated mass?
a) 4.0%
b) 4.2%
c) -4.0%
d) -4.2%
16. The number of electrons in a sulfur atom is:
a. 32.06
b. 32
c. 16
d. 48
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Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.)
17. The sum of the protons and neutrons in a
lithium atom is:
a. 12
b. 13
c. 9
d. 7
18. The electron was discovered by
a. Bohr
b. Thompson
c. Maxwell
d. Dalton
19. The number of neutrons in a fluorine atom is:
a. 19
b. 10
c. 9
d. 18.998
20. The net charge on any atom is:
a. positive
b. negative
c. neutral
d. depends on the element
21. The sum of the protons and neutrons in an atom
equals the:
a. atomic number
b. number of electrons
c. atomic mass
d. mass number
22. Which of these statements is false?
a. Electrons have a negative charge.
b. Electrons have a mass of 1 amu.
c. The nucleus of an atom is positively
charged.
d. The neutron is found in the nucleus of an
atom.
23. An atom with atomic number 48 and mass
number 120 contains:
a. 48 protons, 48 electrons, and 72 neutrons
b. 72 protons, 48 electrons, and 48 neutrons
c. 120 protons, 48 electrons, and 72 neutrons
d. 72 protons, 72 electrons, and 48 neutrons
24. An element which has a mass number of 23 and
has 13 neutrons is the element:
a. Lithium
b. Potassium
c. Magnesium
d. Sodium
d.
Becquerel
26. Rutherford's alpha scattering experiment showed that
the charge on the nucleus of an atom must be:
a. positive
b. neutral
c. negative
d. none of the above
27. The nucleus of the atom has
a. a high density
b. a low density
c. a negative charge
d. no charge
28. An ion always contains
a. unequal number of protons and electrons
b. equal number of protons and electrons
c. unequal number of protons and neutrons
d. equal number of protons and neutrons
29. The whole number that is closest to the atomic mass of
an atom is the
a. atomic number
b. mass number
c. Avogadro's number
d. number of neutrons
30. Which subatomic particle did Thompson include in his
"plumb pudding model"?
a. protons
b. neutrons
c. electrons
d. none of the above
31. Which of the following types of reactions results in a
single product?
a. combination
b. decomposition
c. single replacement
a. double replacement
32. In the reaction 2KClO3  2KCl + 3O2 oxygen is a ___.
a. reactant
b. product
c. coefficient
d. subscript
33. Supersaturated solutions are characterized by
a. being super hot
b. having great stability
c. having a larger amount of solute than can be
dissolved
d. being able to exist at super-low temperatures
25. The experiment that revealed the charge of the
electron involved the use of
a.
gold foil
b.
Rutherford
c.
the Cathode ray
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Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.)
34. In the reaction: N2(g) + 3H2(g) 2NH3(g) + heat
a.
b.
c.
d.
the reaction is both endothermic and exothermic
the reaction is endothermic
the reaction is exothermic
the reaction is neither endothermic or
exothermic
35. How many oxygen atoms are in Al2(SO4)3?
a. 3 atoms of O
b. 4 atoms of O
c. 7 atoms of O
d. d. 12 atoms of O
36. The atomic mass of an element:
a. depends on the number of isotopes of that
element.
b. depends on the mass of each isotope of that
element.
c. depends on the relative abundance of isotopes
of that element.
d. all of the above.
39. The weighted average mass of all the atoms
(isotopes) of an element is the:
a. atomic mass
b. atomic number
c. electron number
d. neutron number
40.
The percentages of isotopes found in a sample
of an element is given below.
22% Carbon-14 with a mass of 14 amu
78% Carbon-12 with a mass of 12 amu
The correct method for finding the average atomic mass
of the sample would be:
a. (22) (14) - (78) (12)
b. (0.22) (14) - (0.78) (12)
c. (0.22) (14) + (0.78) (12)
d. (22) (14) + (78) (12)
41.
How do the isotopes hydrogen-2 and hydrogen-3
differ?
a. Hydrogen-3 has one more electron than
hydrogen-2.
b. Hydrogen-3 has two neutrons.
c. Hydrogen-3 has three protons.
d. Hydrogen-2 has no protons.
42. How many molecules are in 4.50 moles of H2O?
a. 4.50
b. 2.71 x 1024
c. 6.02 x 1023
d. 3.00
43. How many moles are in 8.5 x 10 25 molecules of CO?
a. 1.4 x 102
b. 7.1 x 10-3
c. 5.1 x 1049
d. 8.5 x 1025
44. What is the molar mass of CO2?
a. 36.0 g
b. 22.0 g
c. 44.0 g
d. 6.02 x 1023 g
45. The sum of the atomic masses of all the atoms in a
compound is called the ___.
a. molar mass
b. empirical formula
c. molar volume
d. percentage composition
46. The number of atoms in one mole of an element is
equal to___.
a. a measure
b. a gram
c. a formula unit
d. Avogadro's number
47. It is possible to convert moles to particles by:
a. multiplying by 6.02 x1023
b. dividing by 6.02 x1023
c. multiplying by the molar mass
d. dividing by the molar mass
48. How many molecules of sulfur dioxide are present in
1.60 moles of sulfur dioxide?
a. 9.63 x 1023
b. 102.1 x101
c. 7.62 x 101
d. 3.76 x 1023
49. Find the number of moles in 3.30 g of (NH 4)2SO4
a. 132.1
b. 40.0
c. 0.279
d. 0.0250
50. Which contains more atoms?
a. 1.00 mole H2O2
b. 1.00 mol C2H6
c. 1.00 mol CO
d. 1.00 mol K
51. An element with seven valence electrons would likely be:
a. an alkaline earth metal
b. an alkali metal
c. a noble gas
d. a halogen
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Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.)
52. The most stable atoms are those of the
a. metals
b. metalloids
c. noble gases
d. nonmetals
53. The ion with a charge of +1 and the same
electron configuration as argon is
a. potassium
b. sodium
c. neon
d. magnesium
54. The tendency to lose electrons
______________ as we move across a period
on the periodic table
a. increases
b. remains the same
c. decreases
d. no trend exists
55. Generally, members of a ____________ have
the same number of valence electrons.
a. period
b. series
c. row
d. family
56. An element which is considered to be a
metalloid is:
a. Boron
b. Calcium
c. Oxygen
d. Sodium
57. The element Iodine is a
a. period 5 alkali metal
b. period 4 halogen
c. period 5 halogen
d. period 5 transition metal
58. Sodium and Potassium have similar properties
because they have the same ___.
a. atomic radius
b. number of valence electrons
c. ionization energy
d. electronegativity
59. The maximum number of valence electrons in
an atom is ___.
a. 2
b. 4
c. 8
d. 12
60. The likeliest charge an atom with 2 valence electrons
would develop is
a. 2+
b. 6+
c. 2d. 661. The likeliest charge an atom with 6 valence electrons
will develop is
a. 2+
b. 6+
c. 2d. 662. The likeliest charge of an ion of the element Bromine is
a. -1
b. +1
c. +2
d. -2
63. Metals tend to __.
a. gain electrons and become positively charged
cations
b. lose electrons and become negatively charged
cations
c. lose electrons and become positively charged
cations
d. gain electrons and become negatively charged
anions.
64. The basis of the ionic bond is the ___.
a. sharing of an electron pair
b. electrical attraction between oppositely charged
ions.
c. absorption of energy
d. absorption of water into their solid structures
65. Elements tend to gain or lose electrons in order to
acquire the electron configuration of a
a. halogen
b. transition metal
c. noble gas
d. nonmetal
69. When sodium combines with chlorine to form sodium
chloride, the sodium attains the electron
configuration of
a. helium
b. neon
c. argon
d. lithium
70. The type of chemical bonding in which electron pairs
are shared is
a. ionic
b. covalent
c. metallic
d. none of the above
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Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.)
71. Which of the following correctly matches the
names and formulas of the compounds?
a. AlCl3 . aluminum trichloride and N2O4 ,
nitrogen oxide
b. AlCl3 , aluminum trichloride and N2O4
dinitrogen tetroxide
c. AlCl3, aluminum chloride and N2O4 nitrogen
oxide
d. AlCl3 , aluminum chloride and N2O4 dinitrogen
tetroxide
d.
balanced
74. Choose the correct set of coefficients needed to
balance the following equation:
PCl5  PCl3 + Cl2
a. 2,2,1
b. 1,2,1
c. 1,1,2
d. balanced
75. Choose the correct set of coefficients needed to
balance the following equation:
PbCl2 + Li2SO4  LiCl + PbSO4
a. 1,1,2,1
b. 1,2,2,1
c. 2,2,1,1
d. balanced
73. Choose the correct set of coefficients needed to
balance the following equation:
H2 + O2 ---> H2O
a. 2,1,1
b. 1,1,2
c. 2,1,2
Short Answer:
76. A copper penny has a mass of 3.1 g and a volume of 0.35cm3. What is the density of copper?
8.86 g/cm3
77. A liquid has a density of 4.8 g/l. What is the mass of a 2 liter sample?
9.6 g
78. What is the volume of a substance that has a mass of 80 g and a density of 10 g/cm3?
10 cm3
79. Calculate the following quantities:
a. 1,100 cm
=
1.1 m
b. 1 m
=
1000 mm
c. 10 m
=
1000 cm
d. 2.5 km
=
2500 m.
e. 4.05 kg
=
4050 g
80. Indicate the number of significant figures in
each of the following:
a. 12600
3
b. 0.09
1
c. 2001
4
d. 0.00500100
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83. Which of the following are empirical formulas and
which are molecular formulas?
81. The accepted value or true value for the density
of lead (Pb) is 11.35 g/ml. Your experimental
value or observed value found during a class lab
is 9.65 g/mL
84. Find the empirical formula of each compound from its
% composition.
a. 72.4 % Fe and 27.6% O
___________Fe
What is the error of your measurement?
b. 94.1% O and 5.9%
___________Fe
-1.7
What is the percent error of your measurement? (-)
15%
82. Name the two temperature scales used in science?
Give the freezing pt., and boiling pt. of water for
each of them.
Celcius
0oC / 100oC
Kelvin
a. CH4N
empirical
b. NaO
empirical
c. C6H3O3 molecular
d. H2O2
molecular
e. Na2SO3 empirical
f. C6H10O4 molecular
85. If given the empirical formula and molar mass for a
compound, calculate the compound's molecular
formula?
a. CH2O , mass = 90 g/mol
C3H6O3
b. C3H5O2 mass = 146 g/mol
273 / 373
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C6H10O4
Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.)
86. Find the missing density, mass, or volume of
the following:
a. The mass of a substance is 45.6 g and the volume
is 15 cm3:
Density = ________
b. The volume of a substance is 2.9 ml its density is
6 g/ml:
Mass = ________
95. Calculate the % composition of each element by mass of:
Propane C3H8 _______________ / _______________
Water
______________ / _______________
96. Which of the following are isotopes?
a. 2412X
b. 2713X
c. 2813X
d. 2411X
97. Describe the major parts of Dalton's Atomic Theory:
c. The density of a substance is 7.8 g/cm3 and the
mass is 125 g:
Volume = _______
(Hint: D/1 = M/V (Given any two of the numbers; D, M or V,
you can cross multiply and divide to find what’s missing)
87. If you have 6.7 L of O2 at STP, how many
moles do you have __________________
88. What is the molar mass of Sn3(PO4)2?
_________
98. Balance the following chemical equations:
3 CO + _____ Fe2O3  2 Fe + 3 CO2
3 Zn(OH)2 + 2 H3PO4 ____ Zn3(PO4)2 + 6 H2O
89. How many moles are in 137.5 g of Mn?
________
90. What is the mass of 2 moles of C2H6?
_________
_____ H3PO4 + 3 KOH  _____ K3PO4 + 3 H2O
99. Name and describe the 6 types of chemical reactions.
91. What are the correct formulas or names for the
following compounds?
a. potassium sulfate
K2SO4
b. calcium phosphate
Ca3(PO4)2
c. disulfur heptoxide
S2O7
d. trinitrogen pentahydride N3H5
92. List the diatomic molecules:
a. _______ b. _______ c. _______ d.
_______
100. Describe difference between homogenous,
heterogeneous mixtures, and pure substance:
e. _______ f. _______ g. _______
93. List the names & formulas of the 5 common
acids:
a. ____________________- ______________
b. ____________________- ______________
c. ____________________- ______________
d. ____________________- ______________
e. ____________________- ______________
94. What is the oxidation (nuclear) charge of each
substance (ion) given?
Al __________ Ca __________ O __________
S __________ K __________ N __________
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Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.)
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