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Chemistry I Fall Exam Review
Unit 1: Classification of Matter
1. In order to separate water into the elements of hydrogen and oxygen, a chemical change must take place. This
suggests that water is _____
A. an element.
B. a mixture.
C. a pure substance.
D. heterogeneous.
2.
The state of matter in which a material has definite shape and definite volume is _____
A. liquid.
B. vaporous.
C. gaseous.
D. solid.
3.
A chemical property of oxygen is that is _____
A. has a density of 1.43 g/L.
B. boils at -183°C.
C. is the most abundant element in the earth's crust. D. combines with hydrogen to form water.
4.
The state of matter in which a material has a definite volume but no definite shape is _____
A. gaseous state.
B. frozen state.
C. liquid state.
D. solid state.
5.
Mixtures that are homogeneous are _____
A. pure substances.
C. solutions.
B. composed of regions called phases.
D. not uniform in composition.
6.
All of the following are chemical changes except _____
A. melting of copper.
B. formation of table salt from sodium and chlorine.
C. decaying of plants
D. rusting of iron.
7.
An example of a homogeneous substance is _____
A. pizza.
B. water.
C. vegetable soup.
D. concrete.
8.
When two or more kinds of matter are combined and retain their own characteristic properties within a sample,
that sample is called _____
A. a mineral.
B. a gas.
C. an ore.
D. a mixture.
9.
The state of matter in which material has neither a definite shape nor a definite volume is the _____
A. solid state.
B. elemental state.
C. gaseous state.
D. liquid state.
10. All of the following are usual evidence of a chemical reaction with the exception of _____
A. a solid dissolving.
B. production of light.
C. production of a precipitate.
D. evolution of heat.
11. An experiment that determines the density of water is investigating _____
A. a chemical property. B. a chemical change. C. the potential energy of water. D. a physical property.
Unit 2: Scientific Measurements
12. What is the density of 37.72 g of matter whose volume is 6.80 cm3?
A. 256.4 g/cm3
B. 30.92 g/cm3
C. 0.18g/cm3
D. 5.55 g/cm3
13. Three samples of 0.12 g, 1.8 g, and 0.562 g are weighed and placed together. The combined weight of all three
samples to correct significant figures should be recorded as _____
A. 2.5 g.
B. 2.4 g.
C. 2.482 g.
D. 2.48 g.
14. The sides of a rectangular piece of plywood are 3.54 cm and 4.85 cm. The calculated area is 17.1690 cm 2.
With the correct number of significant figures, the result is expressed as _____
A. 17.169 cm2.
B. 17.17 cm2.
C. 17.1 cm2.
D. 17.2 cm2.
15. The measurement that has only non-significant zeros is _____
A. 0.0037 mL.
B. 60.0 mL
C. 506 mL
D. 400. mL
16. The quantity that has been measured to three significant figures is _____
A. 0.202 g.
B. 3.065 g.
C. 0.052 g.
D. 5000 g.
17. The product of 1.6 cm and 2.4 cm must have how many significant figures?
A. 1
B. 4
C. 2
D. 3
18. The number of significant figures in the measured value of 0.00320 is _____
A. 2.
B. 6.
C. 3.
D. 5.
19. 0.05 cm is the same as _____
A. 0.05 m.
B. 0.005 mm.
D. 0.5 mm.
C. 0.00005 m.
20. The density of sugar is 1.59 g/cm3. A sample with a mass of 4.0 grams would have a volume of _____
A. 6.36 cm3.
B. 2.5 g/cm3.
C. 2.5 cm3.
D. 0.39 cm3.
21. 1.06 L of water is equivalent to _____
A. 1060 mL.
B. 106 mL.
C. 10.6 mL
D. 0.00106 mL.
22. The number of significant figures in the measurement 210 cm is _____
A. 2.
B. 1.
C. 3.
D. 4.
Unit 3: Atomic Structure
23. Isotopes of an element contain different numbers of _____
A. nuclides.
B. protons.
C. neutrons.
D. electrons.
24. Zinc-66 has _____
A. 96 neutrons.
D. 36 neutrons.
B. 30 neutrons.
C. 33 neutrons.
25. A chlorine atom with a mass number of 35 has _____
A. 17 protons, 17 electrons, and 18 neutrons.
B. 17 protons, 17 electrons, and 52 neutrons.
C. 35 protons, 35 electrons, and 17 neutrons.
D. 18 protons, 18 electrons, and 17 neutrons.
26. The nucleus of an atom has all of the following characteristics EXCEPT that it _____
A. contains nearly all of the atom's mass.
B. is positively charged.
C. is very dense.
D. contains nearly all of the atom’s volume.
27. Since any element used in the cathode produced electrons, it was concluded that _____
A. atoms carried a negative charge.
B. only metals contained electrons.
C. all atoms contain electrons.
D. atoms were indivisible.
28. Element X has three isotopes, with the following atomic masses and abundances: X 1 - 3.98 amu, 92.2%; X2 2.99 amu, 6.7%; and X3 - 5.97 amu, 1.1%. Determine the average atomic mass of element X.
A. 3.98 amu
B. 4.15 amu
C. 4.02 amu
D. 3.94 amu
29. An atom is electrically neutral because _____
A. neutrons balance the protons and electrons.
C. nuclear forces equalize the charges.
B. the number of protons and neutrons is equal.
D. the number of protons and electrons is equal.
30. The mass of a neutron is _____
A. double that of an electron.
C. about the same as that of a proton.
B. about the same as that of an electron.
D. double that of a proton.
31. The scientist who discovered the existence of the electron by working with cathode-ray tubes was _____
A. Thomson
B. Bohr
C. Millikan
D. Rutherford
32. Since most particles fired at metal foil passed through the foil, Rutherford concluded that _____
A. atoms were indivisible.
B. atoms contain no charged particles.
C. atoms are mostly empty space.
D. electrons form the nucleus.
33. One part of Dalton's atomic theory that has been modified is the idea that _____
A. atoms cannot be subdivided.
B. all matter is composed of atoms.
C. atoms of different elements have different properties and masses.
D. atoms can combine in chemical reactions.
Unit 4: Electrons in Atoms
34. In the correct Lewis-dot notation for sulfur, the symbol S is surrounded by _____
A. two pairs of dots and two single dots.
B. four single dots.
C. two pairs of dots.
D. three pairs of dots.
35. How many quantum numbers are used to completely describe the energy state of an electron in an atom?
A. 3
B. 2
C. 1
D. 4
36 The electron configuration notation for aluminum is _____
A. ls2 2s2 2p9.
B. ls2 2s2 2p6 3s2 2d1.
C. ls2 2s2 2p3 3s2 3p3 3d1.
D. ls2 2s2 2p6 3s2 3p1.
37. For the p sublevel, the number of orbitals is _____
A. 1.
B. 3.
D. 5.
C. 7.
38. The element with the electron configuration notation ls2 2s2 2p6 3s2 3p2 is _____
A. magnesium.
B. sulfur.
C. silicon.
D. carbon.
39. For an electron in an atom to change from ground state to excited state, _____
A. radiation must be emmited.
B. energy must be absorbed.
C. transitions from higher energy levels to lower energy levels must occur.
D. energy must be released.
40. A three-dimensional region around a nucleus in which a particular electron may be found is called a(n) _____
A. orbit.
B. orbital.
C. electron path.
D. spectral line.
4l. How many elements are in the second period that contain filled or partially filled sand p sublevels?
A. 8
B. 32
C. 2
D. l8
42. Which color of visible light has the shortest wavelength?
A. yellow
B. violet
C. red
D. blue
43. Complete the electron configuration for bromine: ls2 2s2 2p6 _____
A. 3s2 3p6 4s2 4p5
B. 2d10 3s2 3p6 4s2 4p5. C. 3s2 3p6 4s2 4p5 4d10.
D. 3s2 3p6 4s2 3d10 4p5.
44. The number of orbitals for the d sublevel is _____
A. 5
B. 1
D. 3
C. 7
Unit 5: The Periodic Table
45. The atomic radius generally increases with atomic number in a particular group of elements. The dominant
factor that determines this variation is the _____
A. addition of energy levels.
B. increase in the number of neutrons.
C. formation of anew octet.
D. increase in nuclear charge.
46. One-half the distance between the nuclei of identical atoms combined in an element is called the
A. atomic diameter.
B. electron cloud.
C. atomic radius.
D. atomic volume.
47. Mendeleev did not always list elements in his periodic table in order of increasing atomic mass, so that he could
group together elements with similar _____
A. properties.
B. atomic numbers.
C. colors.
D. densities.
48. The periodic law states that the properties of the elements are periodic functions of their atomic numbers. This
means that the _____ determines an element's position in the Periodic Table.
A. number of nucleons. B. number of neutrons. C. number of protons.
D. mass number.
49. The energy required to remove an electron from an atom is the atom's _____
A. ionization energy.
B. electron affinity.
C. electron energy.
D. electronegativity.
50. As you move from left to right across the second period of the Periodic Table, _____
A. The ionization energy increases.
B. The atomic mass decreases.
C. The atomic radii increase.
D. The electronegativity decreases.
51. The vertical columns of the periodic table are called _____
A. periods.
B. rows.
C. transitions.
D. groups.
52. Which group in the Periodic Table would have the lowest first ionization energies?
A. Noble gases
B. Alkali Metals
C. Halogens
D. Alkaline Earth Metals
53. The element that has the greatest electronegativity is _____
A. chlorine
B. fluorine.
C. sodium.
D. oxygen.
Unit 6: Chemical Bonding
54. The substance whose Lewis structure indicates covalent bonds is _____
A. CH2Cl2
B. NH3.
C. H2O
D. CCl4.
55. Use VSEPR theory to predict the shape of hydrogen sulfide, H 2S.
A. Bent
B. Linear
C. Tetrahedral
D. Trigonal Pyramidal
56. The phosphate ion, PO4-3, contains how many extra electrons in its Lewis structure?
A. 2
B. 4
C. 0
D. 3
57. Use VSEPR theory to predict the shape of hydrogen cyanide, HCN.
A. Linear
B. Tetrahedral
C. Trigonal Pyramidal
D. Bent
58. Use VSEPR theory to predict the shape of methane, CH4.
A. Tetrahedral
B. Linear
C. Trigonal Planar
D. Bent
59. If two covalently bonded atoms are identical, the bond is identified as _____
A. nonionic.
B. nonpolar covalent.
C. coordinate covalent.
D. polar covalent.
60. A chemical bond resulting from the sharing of electrons between two atoms is called a(n) _____
A. Lewis structure.
B. ionic bond.
C. orbital bond.
D. covalent bond.
61. The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as _
A. ions.
B. electron clouds.
C. d electrons.
D. valence electrons.
62. In many compounds, atoms of main-group elements form bonds so that the number of electrons in the
outermost energy levels of each atom is _____
A. 2.
B. 8.
C. 6.
D. 10.
63. How many double bonds are in the Lewis structure for hydrogen fluoride, HF.
A. One
B. None
C. Three
D. Two
64. A chemical bond resulting from electrostatic attraction between positive and negative ions is called a(n) _____
A. charged bond.
B. dipole bond.
C. ionic bond.
D. covalent bond.
Unit 7: Chemical Names and Formulas
65. What is the formula for phosphoric acid?
A. H3PO4
B. H3P
C. HPO4
D. H3PO3
66. The name of the compound Cr 2(SO4)3 is _____
A. chromium sulfate
B. chromium (II) sulfate C. chromium (III) sulfate D. chromium sulfur oxide
67. In which of these compounds would the name include the Greek numerical prefixes di- and tri-?
A. Ca3(PO4)2
B. N2O3
C. Fe2O3
D. Al2S3
68. When lead (II) and chromate ions combine, the compound formed has the formula _____
A. Pb2(CrO4)3.
B. PbCrO4.
C. Pb(CrO4)2.
D. Pb2CrO4.
69. What is the formula for silicon dioxide?
A. S2O
B. SiO2
C. Si2O
D. SO2
70. Name the acid with the formula H2SO4.
A. sulfuric acid
B. dyhydrous acid
C. hydrosulfuric acid
D. sulfurous acid
71. Name the compound CF4.
A. carbon fluoride
B. calcium fluoride
C. carbon tetrafluoride
D. monocarbon tetrafluoride
72. Which of the following represents a binary compound?
A. N2O5
B. HClO
C. NaOH
D. KNO3
73. The correct name for the compound that has the formula N2O3 is _____
A. nitrogen trioxide.
B. trinitrogen dioxide.
C. dinitrogen oxide.
D. dinitrogen trioxide.
74. The formula for barium hydroxide is _____
A. BaOH2.
B. BaOH.
D. Ba(OH)2.
C. Ba2OH.
Unit 8: Molar Relationships
75. A compound is 82.4% N and 17.6% H. What is the simplest formula for this compound?
A. N2H
B. NH3
C. NH2
D. NH
76. The formula mass of SiO2 is 60.07 g/mole. How many moles of SiO2 are present in 60.07 g?
A. 60.07 mole
B. 20.02 mole
C. 1.000 mole
D. 100.0 mole
77. What is the mass (in grams) of 4.72 x 1023 molecules CS2?
A. 97.1 g
B. 0.0168 g
C. 59.7 g
D. 0.0102 g
78. Which of the following is NOT a true statement concerning empirical and molecular formulas?
A. The molecular formula of a compound can be some whole-number multiple of its empirical formula.
B. The empirical formula of a compound can be triple it molecular formula.
C. Several compounds can have the same empirical formula, but have different molecular formulas.
D. The molecular formula of a compound can be the same as its empirical formula.
79. Find the volume of 0.250 moles of oxygen gas at STP.
A. 5.60 L
B. 0.0112 L
C. 1.51 x 1023 L
D. 8.00 L
80. A compound analyzed as 80% carbon and 20% hydrogen is found to have a formula mass of 30 g/mole. Its
molecular formula is _____
A. CH3.
B. CH2O.
C. C2H6.
D. C2H4.
81. How many molecules are present in 4.50 L of carbon dioxide gas?
A. 1.21 x 1023 molecules
B. 1.67 x 10-22 molecules
25
C. 6.07 x 10 molecules
D. 3.33 x 10-25 molecules
82. Express 2.50 g NaOH in formula units.
A. 6. 2 x 1025
B. 1.66 x 10-22
C. 3.76 x 1022
D. 1.04 x 10-25
83. All of the following are empirical formulas EXCEPT _____
A. Sn3(PO4)2.
B. N2O4.
C. C6H5Cl.
D. Na2SO4.
84. What is the percent composition of CO?
A. 12 % C, 16 % O
B. 50 % C, 50 % O
D. 43 % C, 57 % O
C. 25 % C, 75 % C
85. The chemical formula of aspirin is C9H8O4. What is the mass of 0.40 moles of aspirin?
A. 10.8 g
B. 160 g
C. 45 g
D. 72 g
86. How many moles of SO3 are in 2.4 x 1024 molecules of SO3?
A. 4.0
B. 3.4 x 1022
C. 2.9 x 10-3
D. 0.25
87. A 75.0 g sample of a compound known to contain only chromium and chlorine was analyzed. The sample was
found to contain 51.3 g chromium and the remainder was chlorine. What is the percent chlorine in the sample?
A. 31.6 %
B. 51.3 %
C. 23.7 %
D. 68.4 %
88. Find the density of sulfur dioxide gas at STP.
A. 1430 g/L
B. 86.46 g/L
C. 2.86 g/L
D. 0.350 g/L
89. The formula mass for copper (II) chloride is _____
A. 134.5 g/mole
B. 99.0 g/mole
C. l62.5 g/mole
D. 98.0 g/mole
Unit 9: Chemical Reactions
90. When you complete and balance the reaction between barium oxide and water, what is the product including the
coefficient?
A. BaH2
B. Ba(OH)2
C. 2BaOH
D. 2Ba(OH)3
91. What is the coefficient of aluminum when the following equation is properly balanced? _Al + _O2  _Al2O3
A. 4
B. 2
C. 6
D. 1
92. What- is the correct reaction for the decomposition of cobalt (III) hydroxide?
A. 2Co(OH)3  2Co(OH)3 + 2Co + 3O2 + 3H2
B. 2Co(OH)3  2Co + 2H2O2
C. 4Co(OH)3  4CoO + O2 + 6H2O
D. 2Co(OH)3  Co2O3 + 3H2O
93. What is the coefficient of bromine after the following equation is properly balanced? _P4 + _Br2  _PBr3
A. 4
B. 6
C. 2
D. 3
94. When copper and sulfur are combined, what product is formed?
A. Cu2S
B. Cu2SO4
C. CuSO4
D. CuS
95. A balanced chemical equation is consistent with the law of _____
A. conservation of mass. B. definite composition. C. multiple proportions.
D. conservation of energy.
96. What is the balanced equation that describes the reaction between calcium carbonate and hydrochloric acid?
A. CaCO3 + 2HCl  CaCl2 + CO2 + H2O
B. CaCO3 + 2HCl  CaCl2 + H2CO3
C. CaCO3 + HCl  CaCl2 + CO2 + H2O
D. CaCO3 + 2HCl  CaCl2 + CO2 + H2
97. What type of reaction does the following equation describe? HgS  Hg + S
A. Synthesis
B. Single Replacement
C. Decomposition
D. Double Replacement
E. Combustion
98. What type of reaction does the following equation describe? 2C6H6 + 15O2  12CO2 + 6H2O
A. Double Replacement B. Single Replacement C. Synthesis
D. Decomposition
E. Combustion
99. Which of the following compounds is insoluble in water?
A. barium sulfate
B. sodium phosphate
C. calcium sulfide
D. silver nitrate
100. What type of reaction does the following equation describe? 2NaBr + I 2  2NaI + Br2
A. Synthesis
B. Combustion
C. Decompositon
D. Single Replacement E. Double Replacement
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