Download Final Exam Study Guide Chapters 1-12

Final Exam Study Guide
Chapters 1-12
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
1. Matter includes all of the following except
a. air.
b. light.
c. smoke.
d. water vapor.
____
2. A physical change occurs when a
a. peach spoils.
b. silver bowl tarnishes.
c. bracelet turns your wrist green.
d. glue gun melts a glue stick.
____
3. Which of the following is an example of a heterogeneous mixture?
a. a gold ring
c. granite
b. seawater
d. sucrose
____
4. The vertical columns on the periodic table are called
a. periods.
c. groups.
b. rows.
d. elements.
____
5. All of the following are steps in the scientific method except
a. observing and recording data.
b. forming a hypothesis.
c. discarding data inconsistent with the hypothesis.
d. developing a model based on experimental results.
____
6. Which of the following observations is qualitative?
a. A chemical reaction was complete in 2.3 seconds.
b. The solid had a mass of 23.4 grams.
c. The pH of a liquid was 5.
d. Salt crystals formed as the liquid evaporated.
____
7. A theory is best described as a
a. series of experimental observations.
b. generalization that explains a body of known facts or phenomena.
c. scientifically proven fact.
d. testable statement.
____
8. Which of these statements does not describe a measurement standard?
a. Measurement standards avoid ambiguity.
b. Measurement standards must be unchanging.
c. A standard can be easily changed to suit the experiment.
d. Confusion is eliminated when the correct measurement is applied.
____
9. The symbol mm represents
a. micrometer.
b. millimeter.
____ 10. The unit m3 measures
a. length.
b. area.
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c. milliliter.
d. meter.
c. volume.
d. time.
Final Exam Study Guide
Chapters 1-12
____ 11. The standard base unit for mass is the
a. gram.
b. cubic centimeter.
c. meter.
d. kilogram.
____ 12. 0.25 g is equivalent to
a. 250 kg.
b. 250 mg.
c. 0.025 mg.
d. 0.025 kg.
____ 13. The number of significant figures in the measured value 0.032 0 g is
a. 2.
c. 4.
b. 3.
d. 5.
____ 14. What is 1.245 633 501  108 rounded to four significant figures?
a. 1246
c. 1.246  108
8
b. 1.2456  10
d. 1.246  104
____ 15. Which of these equations does not describe an inverse proportionality between x and y?
a. xy = k
c. y = k/x
b. x = k/y
d. k = x/y
____ 16. Which of the following is not part of Dalton's atomic theory?
a. Atoms cannot be divided, created, or destroyed.
b. The number of protons in an atom is its atomic number.
c. In chemical reactions, atoms are combined, separated, or rearranged.
d. All matter is composed of extremely small particles called atoms.
____ 17. The most common form of hydrogen has
a. no neutrons.
b. one neutron.
c. two neutrons.
d. three neutrons.
____ 18. How many isotopes of hydrogen are known?
a. 1
b. 2
c. 3
d. 4
____ 19. The nucleus of deuterium contains one proton and
a. two neutrons.
c. no neutrons.
b. one neutron.
d. two electrons.
____ 20. The average atomic mass of an element
a. is the mass of the most abundant isotope.
b. may not equal the mass of any of its isotopes.
c. cannot be calculated.
d. always adds up to 100.
____ 21. Carbon-14 (atomic number 6), the radioactive nuclide used in dating fossils, has
a. 6 neutrons.
c. 10 neutrons.
b. 8 neutrons.
d. 14 neutrons.
____ 22. The number of atoms in 1 mol of carbon is
a. 6.022  1022.
b. 6.022  1023.
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c. 5.022  1022.
d. 5.022  1023.
Final Exam Study Guide
Chapters 1-12
____ 23. Avogadro's number is
a. the maximum number of electrons that all the energy levels can accommodate.
b. the number of protons and neutrons that can fit in the shells of the nucleus.
c. the number of particles in 1 mole of a pure substance.
d. the number of particles in exactly 1 gram of a pure substance.
____ 24. The mass of 2.0 mol of oxygen atoms (atomic mass 16.00 amu) is
a. 16 g.
c. 48 g.
b. 32 g.
d. 64 g.
____ 25. A sample of tin (atomic mass 118.71 amu) contains 3.01  1023 atoms. The mass of the sample is
a. 3.01 g.
c. 72.6 g.
b. 59.3 g.
d. 11 g.
____ 26. The frequency of electromagnetic radiation is measured in waves/second, or
a. nanometers.
c. hertz.
b. quanta.
d. joules.
____ 27. A quantum of electromagnetic energy is called a(n)
a. photon.
c. excited atom.
b. electron.
d. orbital.
____ 28. The wave model of light does not explain
a. the frequency of light.
b. the continuous spectrum.
c. interference.
d. the photoelectric effect.
____ 29. According to the Bohr model of the atom, the single electron of a hydrogen atom circles the nucleus
a. in specific, allowed orbits.
b. in one fixed orbit at all times.
c. at any of an infinite number of distances, depending on its energy.
d. counterclockwise.
____ 30. The region outside the nucleus where an electron can most probably be found is the
a. electron configuration.
c. s sublevel.
b. quantum.
d. electron cloud.
____ 31. According to the quantum theory of an atom, in an orbital
a. an electron's position cannot be known precisely.
b. an electron has no energy.
c. electrons cannot be found.
d. electrons travel around the nucleus on paths of specific radii.
____ 32. The angular momentum quantum number indicates the
a. orientation of an orbital around the nucleus.
b. shape of an orbital.
c. direction of the spin of the electron in its orbital.
d. main energy level of an orbital.
____ 33. The number of possible different orbital shapes for the third energy level is
a. 1.
c. 3.
b. 2.
d. 4.
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Final Exam Study Guide
Chapters 1-12
____ 34. One main energy level can hold 18 electrons. What is n?
a.
c. 6
b. 3
d. 18
____ 35. If the third main energy level contains 15 electrons, how many more could it possibly hold?
a. 0
c. 3
b. 1
d. 17
____ 36. The electron notation for aluminum (atomic number 13) is
a. 1s2 2s2 2p3 3s2 3p3 3d1.
b. 1s2 2s2 2p6 3s2 2d1.
c. 1s2 2s2 2p6 3s2 3p1.
d. 1s2 2s2 2p9.
____ 37. Moseley's work led to the realization that elements with similar properties occurred at regular intervals when
the elements were arranged in order of increasing
a. atomic mass.
c. radioactivity.
b. density.
d. atomic number.
____ 38. The periodic law states that the physical and chemical properties of elements are periodic functions of their
atomic
a. masses.
c. radii.
b. numbers.
d. charges.
____ 39. The periodic law states that
a. no two electrons with the same spin can be found in the same place in an atom.
b. the physical and chemical properties of the elements are functions of their atomic
numbers.
c. electrons exhibit properties of both particles and waves.
d. the chemical properties of elements can be grouped according to periodicity but physical
properties cannot.
____ 40. Period 4 contains 18 elements. How many of these elements have electrons in the d sublevel?
a. 8
c. 16
b. 10
d. 18
____ 41. Neutral atoms with an s2p6 electron configuration in the highest energy level belong to which block of the
periodic table?
a. s block
c. d block
b. p block
d. f block
____ 42. Bromine, atomic number 35, belongs to Group 17. How many electrons does bromine have in its outermost
energy level?
a. 7
c. 18
b. 17
d. 35
____ 43. Magnesium, atomic number 12, has the electron configuration [Ne] 3s2. To what group does magnesium
belong?
a. Group 2
c. Group 5
b. Group 3
d. Group 12
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Final Exam Study Guide
Chapters 1-12
____ 44. The elements in Group 1 are also known as the
a. alkali metals.
c. Period 1 elements.
b. rare-earth series.
d. actinide series.
____ 45. In a row in the periodic table, as the atomic number increases, the atomic radius generally
a. decreases.
c. increases.
b. remains constant.
d. becomes immeasurable.
____ 46. Which is the best reason that the atomic radius generally increases with atomic number in each group of
elements?
a. The nuclear charge increases.
b. The number of neutrons increases.
c. The number of occupied energy levels increases.
d. A new octet forms.
____ 47. Valence electrons are those s and p electrons
a. closest to the nucleus.
b. in the lowest energy level.
c. in the highest energy level.
d. combined with protons.
____ 48. Across a period, ionization energies of d-block elements generally
a. increase.
c. remain constant.
b. decrease.
d. drop to zero.
____ 49. The first electrons to be removed when d-block elements form ions are the
a. d electrons.
c. s electrons.
b. p electrons.
d. f electrons.
____ 50. The chemical formula for an ionic compound represents the
a. number of atoms in each molecule.
b. number of ions in each molecule.
c. ratio of the combined ions present in a sample.
d. total number of ions in the crystal lattice.
____ 51. The chemical formula for water, a covalent compound, is H2O. This formula is an example of a(n)
a. formula unit.
c. ionic formula.
b. Lewis structure.
d. molecular formula.
____ 52. The energy released when 1 mol of an ionic crystalline compound is formed from gaseous ions is called the
a. bond energy.
c. lattice energy.
b. potential energy.
d. energy of crystallization.
____ 53. The forces of attraction between molecules in a molecular compound are
a. stronger than the forces among formula units in ionic bonding.
b. weaker than the forces among formula units in ionic bonding.
c. approximately equal to the forces among formula units in ionic bonding.
d. zero.
____ 54. The Lewis structure for the ammonium ion, NH4, has
a. nonpolar covalent bond.
c. polar covalent bond.
b. ionic bond.
d. metallic bond.
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Final Exam Study Guide
Chapters 1-12
____ 55. In metallic bonds, the mobile electrons surrounding the positive ions are called a(n)
a. Lewis structure.
c. electron cloud.
b. electron sea.
d. dipole.
____ 56. To appear shiny, a material must be able to
a. form crystals.
b. absorb and re-emit light of many wavelengths.
c. absorb light and change it all to energy as heat.
d. change light to electricity.
____ 57. If a material can be shaped or extended by physical pressure, such as hammering, which property does the
material have?
a. conductivity
c. ductility
b. malleability
d. luster
____ 58. Metals are malleable because the metallic bonding
a. holds the layers of ions in rigid positions.
b. maximizes the repulsive forces within the metal.
c. allows one plane of ions to slide past another.
d. is easily broken.
____ 59. Use VSEPR theory to predict the shape of the hydrogen chloride molecule, HCl.
a. tetrahedral
c. bent
b. linear
d. trigonal-planar
____ 60. Use VSEPR theory to predict the shape of the magnesium hydride molecule, MgH2.
a. tetrahedral
c. bent
b. linear
d. octahedral
____ 61. How many atoms of fluorine are present in a molecule of carbon tetrafluoride, CF4?
a. 1
c. 4
b. 2
d. 5
____ 62. The formula for carbon dioxide, CO2, can represent
a. one molecule of carbon dioxide.
b. 1 mol of carbon dioxide molecules.
c. the combination of 1 atom of carbon and 2 atoms of oxygen.
d. all of the above.
____ 63. Name the compound Zn3(PO4)2.
a. zinc potassium oxide
b. trizinc polyoxide
c. zinc phosphate
d. zinc phosphite
____ 64. Name the compound N2O3.
a. dinitrogen oxide
b. nitrogen trioxide
c. nitric oxide
d. dinitrogen trioxide
____ 65. What is the formula for sulfur dichloride?
a. SCl
b. SCl2
c. S2Cl
d. S2Cl2
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Final Exam Study Guide
Chapters 1-12
____ 66. What is the formula for hydrochloric acid?
a. HF
b. HCl
c. HClO
d. H2CO3
____ 67. What is the oxidation number of a pure element?
a. –1
c. +1
b. 0
d. 8
____ 68. What is the oxidation number of oxygen in H2O2?
a. –2
c. +2
b. –1
d. –4
____ 69. Name the compound CCl4 using the Stock system.
a. carbon(IV) chloride
c. carbon chloride
b. carbon tetrachloride
d. carbon hypochlorite
____ 70. Name the compound CO2 using the Stock system.
a. carbon(IV) oxide
c. monocarbon dioxide
b. carbon dioxide
d. carbon oxide
____ 71. The molar mass of LiF is 25.94 g/mol. How many moles of LiF are present in 10.37 g?
a. 0.3998 mol
c. 2.500 mol
b. 1.333 mol
d. 36.32 mol
____ 72. How many Cl– ions are present in 2.00 mol of KCl?
a. 1.20 1024
c. 2.00
b. 6.02 1024
d. 0.5
____ 73. If 0.500 mol of Na+ combines with 0.500 mol of Cl– to form NaCl, how many formula units of NaCl are
present?
a. 3.01 1023
c. 6.02 1024
23
b. 6.02 10
d. 1
____ 74. Of the following molecular formulas for hydrocarbons, which is an empirical formula?
a. CH4
c. C3H6
b. C2H2
d. C4H10
____ 75. A compound's empirical formula is C2H5. If the formula mass is 58 amu, what is the molecular formula?
a. C3H6
c. C5H8
b. C4H10
d. C5H15
____ 76. A compound's empirical formula is NO2. If the formula mass is 92 amu, what is the molecular formula?
a. NO
c. NO4
b. N2O2
d. N2O4
____ 77. How would oxygen be represented in the formula equation for the reaction of methane and oxygen to yield
carbon dioxide and water?
a. oxygen
c. O2
b. O
d. O3
____ 78. In the chemical equation 2Mg(s) + O2(g) ? 2MgO(s),
a. Mg represents the product magnesium.
c. Mg represents the reagent magnesium.
b. the reaction yields magnesium.
d. O2 represents the product oxygen gas.
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Final Exam Study Guide
Chapters 1-12
____ 79. In an equation, the symbol for a substance in water solution is followed by
a. (1).
c. (aq).
b. (g).
d. (s).
____ 80. A chemical formula written over the arrow in a chemical equation signifies
a. a by-product.
c. a catalyst for the reaction.
b. the formation of a gas.
d. an impurity.
____ 81. Which coefficients correctly balance the formula equation CaO + H2O  Ca(OH)2?
a. 2, 1, 2
c. 1, 2, 1
b. 1, 2, 3
d. 1, 1, 1
____ 82. The equation AX  A + X is the general equation for a
a. synthesis reaction.
c. combustion reaction.
b. decomposition reaction.
d. single-displacement reaction.
____ 83. In what kind of reaction do the ions of two compounds exchange places in aqueous solution to form two new
compounds?
a. synthesis reaction
c. decomposition reaction
b. double-displacement reaction
d. combustion reaction
____ 84. The reaction represented by the equation Mg(s) + 2HCl(aq)  H2(g) + MgCl2(aq) is a
a. composition reaction.
c. single-displacement reaction.
b. decomposition reaction.
d. double-displacement reaction.
____ 85. In one type of synthesis reaction, an element combines with oxygen to yield a(n)
a. acid.
c. oxide.
b. hydroxide.
d. metal.
____ 86. When potassium reacts with water, one product formed is
a. hydrogen gas.
c. potassium oxide.
b. oxygen gas.
d. salt.
____ 87. Which reaction can be predicted from the activity series?
a. 2Cl(g)  Cl2(g)
b. HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
c. 2H2O(aq) + 2Na(s)  2NaOH(aq) + H2(g)
d. Cl2(g)  2Cl(g)
____ 88. Which of the following would be investigated in reaction stoichiometry?
a. the masses of hydrogen and oxygen in water
b. the amount of energy released in chemical reactions
c. the mass of potassium required to produce a known mass of potassium chloride
d. the types of bonds that break and form when acids react with metals
____ 89. A balanced chemical equation allows one to determine the
a. mole ratio of any two substances in the reaction.
b. energy released in the reaction.
c. electron configuration of all elements in the reaction.
d. mechanism involved in the reaction.
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Final Exam Study Guide
Chapters 1-12
Use the table below to answer the following questions.
Element
Bromine
Calcium
Carbon
Chlorine
Cobalt
Copper
Fluorine
Hydrogen
Iodine
Iron
Lead
Magnesium
Mercury
Nitrogen
Oxygen
Potassium
Sodium
Sulfur
Symbol
Br
Ca
C
Cl
Co
Cu
F
H
I
Fe
Pb
Mg
Hg
N
O
K
Na
S
Atomic Mass
79.90
40.08
12.01
35.45
58.93
63.55
19.00
1.01
126.90
55.85
207.2
24.30
200.59
14.01
15.00
39.10
22.99
32.01
____ 90. For the reaction represented by the equation Cl2 + 2KBr  2KCl + Br2, how many moles of potassium
chloride are produced from 119 g of potassium bromide?
a. 0.119 mol
c. 0.581 mol
b. 0.236 mol
d. 1.00 mol
____ 91. Which reactant controls the amount of product formed in a chemical reaction?
a. excess reactant
c. composition reactant
b. mole ratio
d. limiting reactant
____ 92. What is the maximum possible amount of product obtained in a chemical reaction?
a. theoretical yield
c. mole ratio
b. percentage yield
d. actual yield
____ 93. For the reaction represented by the equation Mg + 2HCl  H2 + MgCl2, calculate the percentage yield of
magnesium chloride if 100. g of magnesium react with excess hydrochloric acid to yield 330. g of magnesium
chloride.
a. 71.8%
c. 81.6%
b. 74.3%
d. 84.2%
____ 94. Which process can be explained by the kinetic-molecular theory?
a. combustion
c. condensation
b. oxidation
d. displacement reactions
____ 95. An ideal gas is a hypothetical gas
a. not made of particles.
b. that conforms to all of the assumptions of the kinetic theory.
c. whose particles have zero mass.
d. made of motionless particles.
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Final Exam Study Guide
Chapters 1-12
____ 96. According to the kinetic-molecular theory, particles of an ideal gas
a. attract each other but do not collide.
b. repel each other and collide.
c. neither attract nor repel each other but collide.
d. neither attract nor repel each other and do not collide.
____ 97. Under which conditions do real gases most resemble ideal gases?
a. low pressure and low temperature
c. high pressure and high temperature
b. low pressure and high temperature
d. high pressure and low temperature
____ 98. The intermolecular forces between particles in a liquid can involve all of the following except
a. London dispersion forces.
c. dipole-dipole attractions.
b. hydrogen bonding.
d. gravitational forces.
____ 99. Particles within a solid
a. do not move.
b. vibrate about fixed positions.
c. move about freely.
d. exchange positions easily.
____ 100. The compressibility of solids is generally
a. lower than the compressibility of liquids and gases.
b. higher than the compressibility of liquids only.
c. about equal to the compressibility of liquids and gases.
d. higher than the compressibility of gases only.
____ 101. The rate of diffusion in solids is very low because the
a. particles are not free to move about.
b. surfaces of solids usually contact gases.
c. attractive forces are weak.
d. melting points are high.
____ 102. Molecules at the surface of a liquid can enter the vapor phase only if
a. equilibrium has not been reached.
b. the concentration of the vapor is zero.
c. their energy is high enough to overcome the attractive forces in the liquid.
d. condensation is not occurring.
____ 103. If the temperature of a liquid-vapor system at equilibrium increases, the new equilibrium condition will
a. have a lower concentration of vapor.
b. have an increased vapor pressure.
c. not have equal rates of condensation and evaporation.
d. be larger in volume.
____ 104. What is the critical pressure?
a. the pressure at which all substances are solids
b. the pressure at which the attractive forces in matter break down
c. the highest pressure under which a solid can exist
d. the lowest pressure under which a substance can exist as a liquid at the critical temperature
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Final Exam Study Guide
Chapters 1-12
____ 105. At a given temperature, different liquids will have different equilibrium vapor pressures because
a. the energy of the particles is the same for different liquids.
b. diffusion rates differ for the liquids.
c. the attractive forces between the particles differ among liquids.
d. they cannot all be in equilibrium at once.
____ 106. A volatile liquid
a. has strong attractive forces between particles.
b. evaporates readily.
c. has no odor.
d. is ionic.
____ 107. The equilibrium vapor pressure of water is
a. constant at all temperatures.
b. specific for any given temperature.
c. unrelated to temperature.
d. inversely proportional to the temperature.
____ 108. The equilibrium vapor pressure of a liquid increases with increasing temperature because
a. the rate of condensation decreases.
b. the average energy of the particles in the liquid increases.
c. the volume decreases.
d. the boiling point decreases.
____ 109. How does the molar enthalpy of fusion of ice compare with the molar enthalpy of fusion of other solids?
a. It is about the same.
b. It is relatively small.
c. It is relatively large.
d. It is about the same as that of colorless solids.
____ 110. What does the constant bombardment of gas molecules against the inside walls of a container produce?
a. temperature
c. pressure
b. density
d. diffusion
____ 111. Which instrument measures atmospheric pressure?
a. barometer
c. vacuum pump
b. manometer
d. torrometer
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Final Exam Study Guide
Chapters 1-12
Use the table below to answer the following questions.
Water Vapor Pressure
Pressure (mm Hg)
Temperature (C)
0
4.6
5
6.5
10
9.2
15
12.8
20
17.5
25
23.8
30
31.8
35
42.2
40
55.3
50
92.5
____ 112. A sample of gas is collected by water displacement at 600.0 mm Hg and 30C. What is the partial pressure of
the gas?
a. 568.2 mm Hg
c. 630 mm Hg
b. 600.0 mm Hg
d. 631.8 mm Hg
____ 113. At STP, the standard molar volume of a gas of known volume can be used to calculate the
a. number of moles of gas.
c. gram-molecular weight.
b. rate of diffusion.
d. gram-molecular volume.
____ 114. Knowing the mass and volume of a gas at STP allows one to calculate the
a. identity of the gas.
c. condensation point of the gas.
b. molar mass of the gas.
d. rate of diffusion of the gas.
____ 115. When pressure, volume, and temperature are known, the ideal gas law can be used to calculate
a. the chemical formula.
c. molar amount.
b. the ideal gas constant.
d. rate of effusion.
____ 116. What is the pressure exerted by 1.2 mol of a gas with a temperature of 20.C and a volume of 9.5 L?
a. 0.030 atm
c. 3.0 atm
b. 1.0 atm
d. 30. atm
____ 117. Which is an example of effusion?
a. air slowly escaping from a pinhole in a tire
b. the aroma of a cooling pie spreading across a room
c. helium dispersing into a room after a balloon pops
d. oxygen and gasoline fumes mixing in an automobile carburetor
____ 118. Which of the following has components in a nonuniform arrangement?
a. homogeneous mixture
c. salt water
b. solution
d. heterogeneous mixture
____ 119. Colloids
a. can be separated by filtering.
b. settle out when allowed to stand.
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c. scatter light.
d. contain particles larger than 1000 nm.
Final Exam Study Guide
Chapters 1-12
____ 120. The Tyndall effect is used to distinguish between
a. liquids and gases.
c. electrolytes and nonelectrolytes.
b. solutions and colloids.
d. solvents and solutes.
____ 121. If a solution is not agitated while it is being made, dissolved solute tends to
a. mix uniformly.
c. build up in the solvent near the solute.
b. build up in the solvent far from the solute. d. raise the temperature of the solvent.
____ 122. In the expression "like dissolves like," the word like refers to similarity in molecular
a. mass.
c. energy.
b. size.
d. polarity.
____ 123. Pressure has the greatest effect on the solubility of
a. solids in liquids.
c. gases in gases.
b. liquids in liquids.
d. gases in liquids.
____ 124. Which of the following expresses concentration?
a. molality
c. moles of solute per liter of solution
b. molarity
d. All of the above
____ 125. What mass of water must be used to make a 1.35 m solution that contains 8.20 mol of NaOH? (molar mass of
NaOH = 40.00 g/mol)
a. 6.07 kg
c. 11.1 kg
b. 7.44 kg
d. 14.5 kg
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Final Exam Study Guide
Chapters 1-12
Problem
126. Determine the mass in grams of 10.0 mol of bromine. The molar mass of bromine is 79.90 g/mol.
127. Determine the number of moles in 100. g of potassium. The molar mass of potassium is 39.10 g/mol.
Use the periodic table below to answer the following questions.
128. Which element has the following electron configuration: [Ar] 4s2 3d10 4p5?
129. Draw a Lewis structure for the nitrate ion,
. Use VSEPR theory to predict its molecular geometry.
130. The molar mass of aluminum is 26.98 g/mol and the molar mass of fluorine is 19.00 g/mol. Calculate the
molar mass of aluminum trifluoride, AlF3.
131. Write a balanced chemical equation for the following reaction: iron plus copper(I) nitrate yields iron(II)
nitrate plus copper.
132. Tell what type of chemical reaction is represented by the following formula equation. Then balance the
equation.
Ga(s) + S(s)
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Ga2S3(s)
Final Exam Study Guide
Chapters 1-12
133. Use the activity series to determine if the following reaction is possible. If it is possible, write the products
and balance the equation. If not, explain.
4Cr(s) + 3O2(g)
134. Use the activity series to determine if the following reaction is possible. If it is possible, write the products
and balance the equation. If not, explain.
2Sb + 6HCl
135. Given the reaction represented by the equation Mg + 2HCl
H2 + MgCl2, determine to two decimal places
the molar masses of all substances involved. Then write the molar masses as conversion factors.
136. What mass in grams of hydrogen gas is produced if 20.0 mol of Zn are added to excess hydrochloric acid
according to the equation Zn(s) +2HCl(aq)  ZnCl2(aq) + H2(g )?
137. How many grams of ammonium sulfate can be produced if 30.0 mol of H2SO4 react with excess NH3
according to the equation 2NH3(aq) + H2SO4(aq)  (NH4)2SO4(aq)?
138. In the reaction represented by the equation 2NH3 + CO2
CO(NH2)2 + H20, 30.7 g of CO(NH2)2 forms per
1.00 mol of CO2 that reacts when NH3 is in excess. What is the percentage yield?
Use the table below to answer the following questions.
Element
Argon
Bromine
Carbon
Chlorine
Fluorine
Helium
Hydrogen
Nitrogen
Oxygen
Atomic Mass
39.948
79.904
12.011
35.453
18.998
4.0026
1.0079
14.007
15.999
139. A gas sprayed from an aerosol can effuse 2.08 times slower than nitrogen diffuses. What is the molar mass of
the unknown gas?
140. A steam vent releases an unknown gas along with steam. This gas travels 1.56 times as slowly as the steam.
What is the molar mass of the unknown gas?
141. A solution contains 85.0 g of NaNO3, and has a volume of 750. mL. Find the molarity of the solution. (molar
mass of NaCl = 58.44 g/mol)
142. What is the molarity of a solution of sucrose, C12H22O11, that contains 125 g of sucrose in 3.50 L of solution?
(molar mass of C12H22O11 = 342.34 g/mol)
143. How many grams of NaOH are required to prepare 200. mL of a 0.450 M solution? (molar mass of NaOH =
40.00 g/mol)
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144. How many grams of NaC2H3O2 are needed to prepare 350. mL of a 2.75 M solution? (molar mass of
NaC2H3O2 = 82.04 g/mol)
145. How many grams of Na2SO4 are needed to prepare 750. mL of a 0.375 M solution? (molar mass of Na2SO4 =
142.05 g/mol)
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Essay
146. a. What can you determine about the atomic structure of an element if you know the atomic number of the
element and mass numbers of its isotopes? What additional information is needed to determine the average
atomic mass of the element?
b. How do you determine the average atomic mass of the element?
147. a. Explain the significance of Avogadro's constant, 6.022  1023.
b. What is the relationship between it and the molar mass of oxygen, 16.00 g/mol?
148. HO and H2O2 are examples of the empirical and molecular formula of a compound, respectively.
Explain the relationship between these two types of formulas.
149. Consider the equation O2(g) + CS2(l)  CO2(g) + SO2(g).
How does the equation violate the law of conservation of mass? How can the equation be rewritten to
conform to the law of conservation of mass?
150. Hot air balloonists seem to defy gravity.
Explain how hot-air balloonists use Charles’s law.
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