Download Spring 2013 Semester Exam Study Guide (Bonding, Nomenclature

Spring 2013 Semester Exam Study Guide
(Bonding, Nomenclature, Equations, Stoichiometry, Gas Laws, and acids and Bases)
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
1. The electrons involved in the formation of a chemical bond are called
a. dipoles.
c. Lewis electrons.
b. s electrons.
d. valence electrons.
____
2. The electrostatic attraction between positively charged nuclei and negatively charged electrons permits two
atoms to be held together by a(n)
a. chemical bond.
c. neutron.
b. London force.
d. ion.
____
3. Atoms naturally move
a. toward high potential energy.
b. toward low potential energy.
c. toward less stability.
d. away from each other.
____
4. As atoms bond with each other, they
a. increase their potential energy, thus creating less-stable arrangements of matter.
b. decrease their potential energy, thus creating less-stable arrangements of matter.
c. increase their potential energy, thus creating more-stable arrangements of matter.
d. decrease their potential energy, thus creating more-stable arrangements of matter.
____
5. If two covalently bonded atoms are identical, the bond is
a. nonpolar covalent.
c. dipole covalent.
b. polar covalent.
d. coordinate covalent.
____
6. If the atoms that share electrons have an unequal attraction for the electrons, the bond is called
a. nonpolar.
c. ionic.
b. polar.
d. dipolar.
____
7. Most chemical bonds are
a. purely ionic.
b. purely covalent.
c. partly ionic and partly covalent.
d. metallic.
____
8. The B—F bond in BF3 (electronegativity for B is 2.0; electronegativity for F is 4.0) is
a. polar covalent.
c. nonpolar covalent.
b. ionic.
d. metallic.
____
9. In the three molecules, O2, HCl, and F2, what atom would have a partial negative charge?
a. oxygen
c. chlorine
b. hydrogen
d. fluorine
____ 10. A neutral group of atoms held together by covalent bonds is a
a. molecular formula.
c. polyatomic ion.
b. chemical formula.
d. molecule.
____ 11. In a molecule of fluorine, the two shared electrons give each fluorine atom how many electron(s) in the outer
energy level?
a. 1
c. 8
b. 2
d. 32
____ 12. What group of elements satisfies the octet rule without forming compounds?
a. halogen
c. alkali metal
b. noble gas
d. alkaline-earth metal
____ 13. In drawing a Lewis structure, each nonmetal atom except hydrogen should be surrounded by
a. 2 electrons.
c. 8 electrons.
b. 4 electrons.
d. 10 electrons.
____ 14. In drawing a Lewis structure, the central atom is generally the
a. atom with the greatest mass.
b. atom with the highest atomic number.
c. atom with the fewest electrons.
d. least electronegative atom.
____ 15. After drawing a Lewis structure, one should
a. determine the number of each type of atom in the molecule.
b. add unshared pairs of electrons around nonmetal atoms.
c. confirm that the total number of valence electrons used equals the number available.
d. determine the electronegativity of each atom.
____ 16. Multiple covalent bonds may occur in atoms that contain carbon, nitrogen, or
a. chlorine.
c. oxygen.
b. hydrogen.
d. helium.
____ 17. The substance whose Lewis structure shows three covalent bonds is
a. H2O.
c. NH3.
b. CH2Cl2.
d. CCl4.
____ 18. What is the correct Lewis structure for hydrogen chloride, HCl?
a. A
b. B
c. C
d. D
____ 19. What is placed between a molecule's resonance structures to indicate resonance?
a. double-headed arrow
c. series of dots
b. single-headed arrow
d. Lewis structure
____ 20. Chemists once theorized that a molecule that contains a single bond and a double bond split its time existing
as one of these two structures. This effect became known as
a. alternation.
c. multiple bonding.
b. resonance.
d. single-double bonding.
____ 21. 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.
____ 22. A formula that shows only the types and numbers of atoms combined in a single molecule is called a(n)
a. molecular formula.
c. Lewis structure.
b. ionic formula.
d. covalent formula.
____ 23. The ions in most ionic compounds are organized into a
a. molecule.
c. polyatomic ion.
b. Lewis structure.
d. crystal.
____ 24. In a crystal, the electrons of adjacent ions
a. repel each other.
b. attract each other.
c. neutralize each other.
d. have no effect on each other.
____ 25. The lattice energy is a measure of the
a. strength of an ionic bond.
b. strength of a metallic bond.
c. strength of a covalent bond.
d. net charge on a crystal.
____ 26. If the lattice energy of compound A is greater than that of compound B,
a. compound A is not an ionic compound.
b. the bonds in compound A are stronger than the bonds in compound B.
c. compound B is probably a gas.
d. compound A has larger crystals than compound B.
____ 27. Compared with ionic compounds, molecular compounds
a. have higher boiling points.
c. have lower melting points.
b. are brittle.
d. are harder.
____ 28. 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.
____ 29. Ionic compounds are brittle because the strong attractive forces
a. allow the layers to shift easily.
b. cause the compound to vaporize easily.
c. keep the surface dull.
d. hold the layers in relatively fixed positions.
____ 30. The properties of both ionic and molecular compounds are related to the
a. lattice energies of the compounds.
b. strengths of attraction between the particles in the compounds.
c. number of covalent bonds each contains.
d. mobile electrons that they contain.
____ 31. How many extra electrons are in the Lewis structure of the phosphate ion, PO43–?
a. 0
c. 3
b. 2
d. 4
____ 32. How many electrons must be shown in the Lewis structure of the hydroxide ion, OH–?
a. 1
c. 9
b. 8
d. 10
____ 33. A chemical bond formed by the attraction between positive ions and surrounding mobile electrons is a(n)
a. nonpolar covalent bond.
c. polar covalent bond.
b. ionic bond.
d. metallic bond.
____ 34. Compared with nonmetals, the number of valence electrons in metals is generally
a. smaller.
c. about the same.
b. greater.
d. almost triple.
____ 35. In metals, the valence electrons
a. are attached to particular positive ions.
b. are shared by all of the atoms.
c. are immobile.
d. form covalent bonds.
____ 36. 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.
____ 37. As light strikes the surface of a metal, the electrons in the electron sea
a. allow the light to pass through.
b. become attached to particular positive ions.
c. fall to lower energy levels.
d. absorb and re-emit the light.
____ 38. 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.
____ 39. Shifting the layers of an ionic crystal causes the crystal to
a. be drawn into a wire.
c. become metallic.
b. shatter.
d. emit light.
____ 40. The concept that electrostatic repulsion between electron pairs surrounding an atom causes these pairs to be
separated as far as possible is the foundation of
a. VSEPR theory.
c. the electron-sea model.
b. the hybridization model.
d. Lewis theory.
____ 41. The strong forces of attraction between the positive and negative regions of molecules are called
a. dipole-dipole forces.
c. lattice forces.
b. London forces.
d. orbital forces.
____ 42. Compared with molecular bonds, the strength of intermolecular forces is
a. weaker.
c. about the same.
b. stronger.
d. too variable to compare.
____ 43. The equal but opposite charges present in the two regions of a polar molecule create a(n)
a. electron sea.
c. crystal lattice.
b. dipole.
d. ionic bond.
____ 44. A polar molecule contains
a. ions.
b. a region of positive charge and a region of negative charge.
c. only London forces.
d. no bonds.
____ 45. A chemical formula includes the symbols of the elements in the compound and subscripts that indicate
a. atomic mass of each element.
b. number of atoms or ions of each element that are combined in the compound.
c. formula mass.
d. charges on the elements or ions.
____ 46. How many atoms of fluorine are present in a molecule of carbon tetrafluoride, CF4?
a. 1
c. 4
b. 2
d. 5
____ 47. 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.
____ 48. What is the formula for zinc fluoride?
a. ZnF
b. ZnF2
c. Zn2F
d. Zn2F3
____ 49. What is the formula for the compound formed by lead(II) ions and chromate ions?
a. PbCrO4
c. Pb2(CrO4)3
b. Pb2CrO4
d. Pb(CrO4)2
____ 50. What is the formula for aluminum sulfate?
a. AlSO4
b. Al2SO4
c. Al2(SO4)3
d. Al(SO4)3
____ 51. What is the formula for barium hydroxide?
a. BaOH
b. BaOH2
c. Ba(OH)2
d. Ba(OH)
____ 52. Name the compound Ni(ClO3)2.
a. nickel(II) chlorate
b. nickel(II) chloride
c. nickel(II) chlorite
d. nickel(II) peroxide
____ 53. Name the compound Zn3(PO4)2.
a. zinc potassium oxide
b. trizinc polyoxide
c. zinc phosphate
d. zinc phosphite
____ 54. Name the compound KClO3.
a. potassium chloride
b. potassium trioxychlorite
c. potassium chlorate
d. hypochlorite
____ 55. Name the compound Fe(NO3)2.
a. iron(II) nitrate
b. iron(II) nitrite
c. iron(III) nitrate
d. iron(III) nitride
____ 56. Name the compound Al2S3.
a. aluminum sulfate
b. aluminum sulfur
c. aluminum(II) sulfate
d. aluminum sulfide
____ 57. Name the compound CF4.
a. calcium fluoride
b. carbon fluoride
c. carbon tetrafluoride
d. monocalcium quadrafluoride
____ 58. Name the compound SiO2.
a. silver oxide
b. silicon oxide
c. silicon dioxide
d. monosilver dioxide
____ 59. Name the compound SO3.
a. sulfur trioxide
b. silver trioxide
c. selenium trioxide
d. sodium trioxide
____ 60. Name the compound N2O3.
a. dinitrogen oxide
b. nitrogen trioxide
c. nitric oxide
d. dinitrogen trioxide
____ 61. What is the formula for silicon dioxide?
a. SO2
b. SiO2
c. Si2O
d. S2O
____ 62. What is the formula for dinitrogen trioxide?
a. Ni2O3
b. NO3
c. N2O6
d. N2O3
____ 63. What is the formula for sulfur dichloride?
a. SCl
b. SCl2
c. S2Cl
d. S2Cl2
____ 64. What is the formula for diphosphorus pentoxide?
a. P2PeO5
c. P2O4
b. PO5
d. P2O5
____ 65. What is the formula for hydrochloric acid?
a. HF
b. HCl
c. HClO
d. H2CO3
____ 66. Name the compound PBr5 using the Stock system.
a. potassium hexabromide
c. phosphorus(V) bromide
b. phosphorus(V) pentabromide
d. phosphoric acid
____ 67. What is the formula mass of ethyl alcohol, C2H5OH?
a. 30.33 amu
c. 45.06 amu
b. 33.27 amu
d. 46.08 amu
____ 68. What is the formula mass of (NH4)2SO4?
a. 114.09 amu
b. 118.34 amu
c. 128.06 amu
d. 132.16 amu
____ 69. The molar mass of MgI2 is
a. the sum of the masses of 1 mol of Mg and 2 mol of I.
b. the sum of the masses of 1 mol of Mg and 1 mol of I.
c. the sum of the masses of 1 atom of Mg and 2 atoms of I.
d. the sum of the masses of 1 atom of Mg and 1 atom of I.
____ 70. The molar mass of CCl4 is 153.81 g/mol. How many grams of CCl4 are needed to have 5.000 mol?
a. 5 g
c. 769.0 g
b. 30.76 g
d. 796.05 g
____ 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. The molar mass of CS2 is 76.15 g/mol. How many grams of CS2 are present in 10.00 mol?
a. 0.13 g
c. 10.00 g
b. 7.614 g
d. 761.5 g
____ 73. The molar mass of NH3 is 17.03 g/mol. How many moles of NH3 are present in 107.1 g?
a. 0.1623 mol
c. 6.289 mol
b. 3.614 mol
d. 107.1 mol
____ 74. What is the mass of 0.240 mol glucose, C6H12O6?
a. 24.0 g
c. 180.16 g
b. 43.2 g
d. 750. g
____ 75. How many Mg2+ ions are found in 1.00 mol of MgO?
a. 3.01 1023
c. 12.04 1023
23
b. 6.02 10
d. 6.02 1025
____ 76. What is the percentage composition of CuCl2?
a. 33% Cu, 66% Cl
c. 65.50% Cu, 34.50% Cl
b. 50% Cu, 50% Cl
d. 47.27% Cu, 52.73% Cl
____ 77. What is the mass percentage of OH– in Ca(OH)2?
a. 45.9%
c. 75%
b. 66.6%
d. 90.1%
____ 78. What is the mass percentage of chlorine in NaCl?
a. 35.45%
c. 60.7%
b. 50%
d. 64.5%
____ 79. A formula that shows the simplest whole-number ratio of the atoms in a compound is the
a. molecular formula.
c. experimental formula.
b. ideal formula.
d. empirical formula.
____ 80. The empirical formula is always the accepted formula for a(n)
a. atom.
c. molecular compound.
b. molecule.
d. ionic compound.
____ 81. The empirical formula for a compound shows the symbols of the elements with subscripts indicating the
a. actual numbers of atoms in a molecule.
b. number of moles of the compound in 100 g.
c. smallest whole-number ratio of the atoms.
d. atomic masses of each element.
____ 82. The empirical formula may not represent the actual composition of a unit of a(n)
a. ionic compound.
c. salt.
b. molecular compound.
d. crystal.
____ 83. To find the molecular formula from the empirical formula, one must determine the compound's
a. density.
c. structural formula.
b. formula mass.
d. crystal lattice.
____ 84. The molecular formula for vitamin C is C6H8O6. What is the empirical formula?
a. CHO
c. C3H4O3
b. CH2O
d. C2H4O2
____ 85. Of the following molecular formulas for hydrocarbons, which is an empirical formula?
a. CH4
c. C3H6
b. C2H2
d. C4H10
____ 86. A compound's empirical formula is N2O5. If the formula mass is 108 amu, what is the molecular formula?
a. N2O5
c. NO3
b. N4O10
d. N2O4
____ 87. A compound's empirical formula is HO. If the formula mass is 34 amu, what is the molecular formula?
a. H2O
c. HO3
b. H2O2
d. H2O3
____ 88. In a chemical reaction
a. the mass of the reactants equals the mass of the products.
b. the mass of the products is greater than the mass of reactants.
c. the number of atoms in the reactants and products must change.
d. energy as heat must be added to the reactants.
____ 89. When a solid produced by a chemical reaction separates from the solution it is called
a. a precipitate.
c. a molecule.
b. a reactant.
d. the mass of the product.
____ 90. The word equation solid carbon + oxygen gas  carbon dioxide gas + energy, represents a chemical reaction
because
a. the reaction releases energy.
b. CO2 has chemical properties that differ from those of C and O.
c. the reaction absorbs energy.
d. CO2 is a gas and carbon is a crystal.
____ 91. In writing a chemical equation that produces hydrogen gas, the correct representation of hydrogen gas is
a. H.
c. H2.
b. 2H.
d. OH.
____ 92. What is the small whole number that appears in front of a formula in a chemical equation?
a. a subscript
c. a ratio
b. a superscript
d. a coefficient
____ 93. To balance a chemical equation, it may be necessary to adjust the
a. coefficients.
c. formulas of the products.
b. subscripts.
d. number of products.
____ 94. According to the law of conservation of mass, the total mass of the reacting substances is
a. always more than the total mass of the products.
b. always less than the total mass of the products.
c. sometimes more and sometimes less than the total mass of the products.
d. always equal to the total mass of the products.
____ 95. A chemical equation is balanced when the
a. coefficients of the reactants equal the coefficients of the products.
b. same number of each kind of atom appears in the reactants and in the products.
c. products and reactants are the same chemicals.
d. subscripts of the reactants equal the subscripts of the products.
____ 96. In the word equation, sodium oxide + water  sodium hydroxide, the formula for
sodium hydroxide is represented by
a. Na2OH.
c. NaO2.
b. NaOH.
d. Na2O.
____ 97. Which word equation represents the reaction that produces water from hydrogen and oxygen?
a. Water is produced from hydrogen and oxygen.
b. Hydrogen plus oxygen yields water.
c. H2 + O2  water.
d. Water can be separated into hydrogen and oxygen.
____ 98. 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
____ 99. Which of the following is a formula equation for the formation of carbon dioxide from carbon and oxygen?
a. Carbon plus oxygen yields carbon dioxide. c. CO2  C + O2
b. C + O2  CO2
d. 2C + O  CO2
____ 100. In an equation, the symbol for a substance in water solution is followed by
a. (1).
c. (aq).
b. (g).
d. (s).
____ 101. 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.
____ 102. When the equation Fe3O4 + Al  Al2O3 + Fe is correctly balanced, what is the coefficient of Fe?
a. 3
c. 6
b. 4
d. 9
____ 103. Which coefficients correctly balance the formula equation
NH4NO2(s) N2(g) + H2O(l)?
a. 1, 2, 2
c. 2, 1, 1
b. 1, 1, 2
d. 2, 2, 2
____ 104. 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
____ 105. The complete balanced equation for the reaction between zinc hydroxide and acetic acid is
a. ZnOH + CH3COOH  ZnCH3COO + H2O.
b. Zn(OH)2 + CH3COOH  Zn + 2CO2 +3H2O.
c. Zn(OH)2 + 2CH3COOH  Zn(CH3COO)2 + 2H2O.
d. Zn(OH)2 + 2CH3COOH  Zn(CH3COO)2 + H2 + O2.
____ 106. What is the balanced equation for the combustion of sulfur?
a. S(s) + O2(g)  SO(g)
b. S(s) + O2(g)  SO2(g)
c. 2S(s) + 3O2(g)  SO3(s)
d. S(s) + 2O2(g)  SO42–(aq)
____ 107. Which equation is not balanced?
a. 2H2 + O2  2H2O
b. 4H2 + 2O2  4H2O
c. H2 + H2 + O2  H2O + H2O
d. 2H2 + O2  H2O
____ 108. In what kind of reaction do two or more substances combine to form a new compound?
a. decomposition reaction
c. double-displacement reaction
b. ionic reaction
d. synthesis reaction
____ 109. In what kind of reaction does one element replace a similar element in a compound?
a. displacement reaction
c. decomposition reaction
b. combustion
d. ionic reaction
____ 110. In what kind of reaction does a single compound produce two or more simpler substances?
a. decomposition reaction
c. single-displacement reaction
b. synthesis reaction
d. ionic reaction
____ 111. 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
____ 112. The reaction represented by the equation 2Mg(s) + O2(g)  2MgO(s) is a
a. synthesis reaction.
c. single-displacement reaction.
b. decomposition reaction.
d. double-displacement reaction.
____ 113. 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.
____ 114. The reaction represented by the equation 2HgO(s)  2Hg(l) + O2(g) is a(n)
a. single-displacement reaction.
c. combustion reaction.
b. synthesis reaction.
d. decomposition reaction.
____ 115. The reaction represented by the equation Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq) is a
a. double-displacement reaction.
c. decomposition reaction.
b. synthesis reaction.
d. combustion reaction.
____ 116. The reaction represented by the equation 2KClO3(s)  2KCl(s) + 3O2(g) is a(n)
a. synthesis reaction.
c. combustion reaction.
b. decomposition reaction.
d. ionic reaction.
____ 117. The reaction represented by the equation Cl2(g) + 2KBr(aq)  2KCl(aq) + Br2(l) is a(n)
a. synthesis reaction.
c. single-displacement reaction.
b. decomposition reaction.
d. combustion reaction.
____ 118. In the equation 2Al(s) + 3Fe(NO3)2(aq)  3Fe(s) + 2Al(NO3)3(aq), iron has been replaced by
a. nitrate.
c. aluminum.
b. water.
d. nitrogen.
____ 119. In a double-displacement reaction, hydrogen chloride and sodium hydroxide react to produce sodium
chloride. Another product is
a. sodium hydride.
c. water.
b. potassium chloride.
d. hydrogen gas.
____ 120. The formulas for the products of the reaction between sodium hydroxide and sulfuric acid are
a. Na2SO4 and H2O.
c. SI4 and Na2O.
b. NaSO4 and H2O.
d. S + O2 and Na.
____ 121. What is the balanced equation when aluminum reacts with copper(II) sulfate?
a. Al + Cu2S  Al2S + Cu
b. 2Al + 3CuSO4  Al2(SO4)3 + 3Cu
c. Al + CuSO4  AlSO4 + Cu
d. 2Al + Cu2SO4  Al2SO4 + 2Cu
____ 122. The ability of an element to react is the element's
a. valence.
c. stability.
b. activity.
d. electronegativity.
____ 123. What is the name of a list of elements arranged according to the ease with which they undergo certain
chemical reactions?
a. reactivity list
c. activity series
b. reaction sequence
d. periodic list
____ 124. An element in the activity series can replace any element
a. in the periodic table.
c. above it on the list.
b. below it on the list.
d. in its group.
____ 125. What can be predicted by using an activity series?
a. whether a certain chemical reaction will occur
b. the amount of energy released by a chemical reaction
c. the electronegativity values of elements
d. the melting points of elements
____ 126. If a certain metal is placed in an ionic solution containing another metal and no reaction occurs, then the metal
originally in the solution is
a. a halogen.
c. not on the activity series.
b. higher on the activity series.
d. unreactive.
____ 127. Predict what happens when calcium metal is added to a solution of magnesium chloride.
a. No reaction occurs.
c. Magnesium calcite forms.
b. Calcium chloride forms.
d. Gaseous calcium is produced.
____ 128. Predict what happens when zinc is added to water.
a. No reaction occurs.
c. Zinc oxide forms.
b. Steam is produced.
d. Hydrogen is released.
____ 129. Predict what happens when lead is added to nitric acid.
a. No reaction occurs.
c. Lead oxide forms.
b. Oxygen is released.
d. Hydrogen is released.
____ 130. Predict what happens when nickel is added to a solution of potassium chloride.
a. No reaction occurs.
c. Potassium nickel chloride forms.
b. Nickel chloride forms.
d. Hydrochloric acid forms.
____ 131. Magnesium bromide (aq) + chlorine (g) yields
a. Mg(s) and BrCl(aq).
c. MgBrCl(aq).
b. MgCl(aq) and Br2(l).
d. MgCl2(aq) and Br2(l).
____ 132. Which reaction does not occur?
a. 2HF(aq) + Cl2(g)  F2(g) + 2HCl(aq)
b. 2Na(s) + ZnF2(aq)  2NaF(aq) + Zn(s)
c. Fe(s) + CuCl2(aq)  FeCl2(aq) + Cu(s)
d. 2HCl(aq) + Mg(s)  MgCl2(aq) + H2(g)
____ 133. Which branch of chemistry deals with the mass relationships of elements in compounds and the mass
relationships among reactants and products in chemical reactions?
a. qualitative analysis
c. chemical kinetics
b. entropy
d. stoichiometry
____ 134. The coefficients in a chemical equation represent the
a. masses, in grams, of all reactants and products.
b. relative numbers of moles of reactants and products.
c. number of atoms in each compound in a reaction.
d. number of valence electrons involved in the reaction.
____ 135. Which equation is not balanced?
a.
b.
c.
d.
____ 136. Which coefficients correctly balance the formula
a. 1,2,2
c. 2,1,1
b. 1,1,2
d. 2,2,2
?
____ 137. When the formula equation
a. 3.
b. 2.
is correctly balanced the coefficient of Fe is number
c. 7.
d. 9.
____ 138. The units of molar mass are
a. g/mol.
b. mol/g.
c. amu/mol.
d. amu/g.
____ 139. In the reaction represented by the equation N2 + 3H2  2NH3, what is the mole ratio of nitrogen to ammonia?
a. 1:1
c. 1:3
b. 1:2
d. 2:3
____ 140. In the reaction represented by the equation C + 2H2  CH4, what is the mole ratio of hydrogen to methane?
a. 1:1
c. 1:2
b. 2:1
d. 2:4
____ 141. In the reaction represented by the equation N2 + 3H2  2NH3, what is the mole ratio of hydrogen to
ammonia?
a. 1:1
c. 3:2
b. 2:1
d. 6:8
____ 142. The Haber process for producing ammonia commercially is represented by the equation N2(g) + 3H2(g) 
2NH3(g). To completely convert 9.0 mol hydrogen gas to ammonia gas, how many moles of nitrogen gas are
required?
a. 1.0 mol
c. 3.0 mol
b. 2.0 mol
d. 6.0 mol
____ 143. For the reaction represented by the equation N2 + 3H2  2NH3, how many moles of nitrogen are required to
produce 18 mol of ammonia?
a. 9.0 mol
c. 27 mol
b. 18 mol
d. 36 mol
____ 144. For the reaction represented by the equation AgNO3 + NaCl  NaNO3 + AgCl, how many moles of silver
chloride, AgCl, are produced from 7.0 mol of silver nitrate AgNO3?
a. 1.0 mol
c. 7.0 mol
b. 2.3 mol
d. 21 mol
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
____ 145. For the reaction represented by the equation SO3 + H2O  H2SO4, how many grams of sulfur trioxide are
required to produce 4.00 mol of sulfuric acid in an excess of water?
a. 80.0 g
c. 240. g
b. 160. g
d. 320. g
____ 146. For the reaction represented by the equation 2Fe + O2  2FeO, how many grams of iron(II) oxide are
produced from 8.00 mol of iron in an excess of oxygen?
a. 71.8 g
c. 712 g
b. 575 g
d. 1310 g
____ 147. For the reaction represented by the equation 2Na + Cl2  2NaCl, how many grams of chlorine gas are
required to react completely with 2.00 mol of sodium?
a. 35.5 g
c. 141.8 g
b. 70.9 g
d. 212.7 g
____ 148. For the reaction represented by the equation 2HNO3 + Mg(OH)2  Mg(NO3)2 + 2H2O, how many grams of
magnesium nitrate are produced from 8.00 mol of nitric acid, HNO3, and an excess of Mg(OH)2?
a. 148 g
c. 593 g
b. 445 g
d. 818 g
____ 149. For the reaction represented by the equation CH4 + 2O2  CO2 + 2H2O, how many moles of carbon dioxide
are produced from the combustion of 100. g of methane?
a. 6.23 mol
c. 12.5 mol
b. 10.8 mol
d. 25 mol
____ 150. For the reaction represented by the equation Pb(NO3)2 + 2KI  PbI2 + 2KNO3, how many moles of lead(II)
iodide are produced from 300. g of potassium iodide and an excess of Pb(NO3)2?
a. 0.904 mol
c. 3.61 mol
b. 1.81 mol
d. 11.0 mol
____ 151. 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
____ 152. For the reaction represented by the equation 3Fe + 4H2O  Fe3O4 + 4H2, how many moles of iron(III) oxide
are produced from 500. g of iron in an excess of H2O?
a. 1.04 mol
c. 8.95 mol
b. 2.98 mol
d. 12.98 mol
____ 153. For the reaction represented by the equation 2KlO3  2KCl + 3O2, how many moles of potassium chlorate
are required to produce 250. g of oxygen?
a. 2.00 mol
c. 4.97 mol
b. 4.32 mol
d. 5.21 mol
____ 154. Ozone, O3, is produced by the reaction represented by the following equation:
What mass of ozone will form from the reaction of 2.0 g of NO2 in a car's exhaust and excess oxygen?
a. 1.1 g O3
c. 2.1 g O3
b. 1.8 g O3
d. 4.2 g O3
____ 155. For the reaction represented by the equation Cl2 + 2KBr  2KCl + Br2, how many grams of potassium
chloride can be produced from 300. g each of chlorine and potassium bromide?
a. 98.7 g
c. 188 g
b. 111 g
d. 451 g
____ 156. For the reaction represented by the equation 2Na + 2H2O  2NaOH + H2, how many grams of hydrogen are
produced if 120. g of sodium and 80. g of water are available?
a. 4.5 g
c. 80. g
b. 45 g
d. 200 g
____ 157. For the reaction represented by the equation 2Na + Cl2  2NaCl, how many grams of sodium chloride can be
produced from 500. g each of sodium and chlorine?
a. 112 g
c. 409 g
b. 319 g
d. 824 g
____ 158. For the reaction represented by the equation SO3 + H2O  H2SO4, how many grams of sulfuric acid can be
produced from 200. g of sulfur trioxide and 100. g of water?
a. 100. g
c. 245 g
b. 200. g
d. 285 g
____ 159. Which reactant controls the amount of product formed in a chemical reaction?
a. excess reactant
c. composition reactant
b. mole ratio
d. limiting reactant
____ 160. A chemical reaction involving substances A and B stops when B is completely used. B is the
a. excess reactant.
c. primary reactant.
b. limiting reactant.
d. primary product.
____ 161. When the limiting reactant in a chemical reaction is completely used, the
a. excess reactants begin combining.
c. reaction speeds up.
b. reaction slows down.
d. reaction stops.
____ 162. After calculating the amount of reactant B required to completely react with A, then comparing that amount
with the amount of B available, one can determine the
a. limiting reactant.
c. energy released in the reaction.
b. rate of the reaction.
d. pathway of the reaction.
____ 163. What is the measured amount of a product obtained from a chemical reaction?
a. mole ratio
c. theoretical yield
b. percentage yield
d. actual yield
____ 164. In most chemical reactions the amount of product obtained is
a. equal to the theoretical yield.
c. more than the theoretical yield.
b. less than the theoretical yield.
d. more than the percentage yield.
____ 165. 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
____ 166. For the reaction represented by the equation SO3 + H2O  H2SO4, calculate the percentage yield if 500. g of
sulfur trioxide react with excess water to produce 575 g of sulfuric acid.
a. 82.7%
c. 91.2%
b. 88.3%
d. 93.9%
____ 167. For the reaction represented by the equation Cl2 + 2KBr  2KCl + Br2, calculate the percentage yield if 200.
g of chlorine react with excess potassium bromide to produce 410. g of bromine.
a. 73.4%
c. 91.0%
b. 82.1%
d. 98.9%
____ 168. For the reaction represented by the equation CH4 + 2O2  2H2O + CO2, calculate the percentage yield of
carbon dioxide if 1000. g of methane react with excess oxygen to produce 2300. g of carbon dioxide.
a. 83.88%
c. 92.76%
b. 89.14%
d. 96.78%
____ 169. 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%
____ 170. Pressure is the force per unit
a. volume.
b. surface area.
c. length.
d. depth.
____ 171. What does the constant bombardment of gas molecules against the inside walls of a container produce?
a. temperature
c. pressure
b. density
d. diffusion
____ 172. Which instrument measures atmospheric pressure?
a. barometer
c. vacuum pump
b. manometer
d. torrometer
____ 173. Standard temperature is exactly
a. 100ºC.
b. 273ºC.
c. 0ºC.
d. 0 K.
____ 174. Standard pressure is exactly
a. 1 atm.
b. 760 atm.
c. 101.325 atm.
d. 101 atm.
____ 175. Who developed the concept that the total pressure of a mixture of gases is the sum of their partial pressures?
a. Charles
c. Kelvin
b. Boyle
d. Dalton
____ 176. The volume of a gas is 400.0 mL when the pressure is 1.00 atm. At the same temperature, what is the pressure
at which the volume of the gas is 2.0 L?
a. 0.5 atm
c. 0.20 atm
b. 5.0 atm
d. 800 atm
____ 177. A sample of oxygen occupies 560. mL when the pressure is 800.00 mm Hg. At constant temperature, what
volume does the gas occupy when the pressure decreases to 700.0 mm Hg?
a. 80.0 mL
c. 600. mL
b. 490. mL
d. 640. mL
____ 178. If the temperature of a fixed quantity of gas decreases and the pressure remains unchanged,
a. its volume increases.
c. its volume decreases.
b. its volume is unchanged.
d. its density decreases.
____ 179. The volume of a gas is 5.0 L when the temperature is 5.0ºC. If the temperature is increased to 10.0ºC without
changing the pressure, what is the new volume?
a. 2.5 L
c. 5.1 L
b. 4.8 L
d. 10.0 L
____ 180. At 7.0ºC, the volume of a gas is 49 mL. At the same pressure, its volume is 74 mL at what temperature?
a. 3.0ºC
c. 120ºC
b. 423ºC
d. 150ºC
____ 181. Why does the air pressure inside the tires of a car increase when the car is driven?
a. Some of the air has leaked out.
b. The air particles collide with the tire after the car is in motion.
c. The air particles inside the tire increase their speed because their temperature rises.
d. The atmosphere compresses the tire.
____ 182. On a cold winter morning when the temperature is –13ºC, the air pressure in an automobile tire is 1.5 atm. If
the volume does not change, what is the pressure after the tire has warmed to 15ºC?
a. –1.5 atm
c. 3.0 atm
b. 1.7 atm
d. 19.5 atm
____ 183. The pressure of a sample of gas at 10.0ºC increases from 700. mm Hg to 900. mm Hg. What is the new
temperature?
a. 0ºC
c. 39.0ºC
b. 364ºC
d. 90.9ºC
____ 184. The volume of a sample of oxygen is 300.0 mL when the pressure is 1.00 atm and the temperature is 27.0ºC.
At what temperature is the volume 1.00 L and the pressure 0.500 atm?
a. 22.0ºC
c. 0.50 K
b. 45.0ºC
d. 227ºC
____ 185. The volume of a gas collected when the temperature is 11.0ºC and the pressure is 710 mm Hg measures 14.8
mL. What is the calculated volume of the gas at 20.0ºC and 740 mm Hg?
a. 7.8 mL
c. 14.6 mL
b. 13.7 mL
d. 15 mL
____ 186. Calculate the approximate volume of a 0.600 mol sample of gas at 15.0°C and a pressure of 1.10 atm.
a. 12.9 L
c. 24.6 L
b. 22.4 L
d. 139 L
____ 187. Calculate the approximate temperature of a 0.50 mol sample of gas at 750 mm Hg and a volume of 12 L.
a. –7ºC
c. 15ºC
b. 11ºC
d. 288ºC
____ 188. What is the pressure exerted by 1.2 mol of a gas with a temperature of 20.ºC and a volume of 9.5 L?
a. 0.030 atm
c. 3.0 atm
b. 1.0 atm
d. 30. atm
____ 189. A sample of gas at 25ºC has a volume of 11 L and exerts a pressure of 660 mm Hg. How many moles of gas
are in the sample?
a. 0.39 mol
c. 9.3 mol
b. 3.9 mol
d. 87 mol
____ 190. If a gas with an odor is released in a room, it can quickly be detected across the room because it
a. diffuses.
c. is compressed.
b. is dense.
d. condenses.
____ 191. 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
____ 192. Acids taste
a. sweet.
b. sour.
c. bitter.
d. salty.
____ 193. Acids generally release H2 gas when they react with
a. nonmetals.
c. active metals.
b. semimetals.
d. inactive metals.
____ 194. Acids react with
a. bases to produce salts and water.
b. salts to produce bases and water.
c. water to produce bases and salts.
d. neither bases, salts, nor water.
____ 195. Aqueous solutions of acids
a. contain only two different elements.
b. carry electricity.
c. have very high boiling points.
d. cannot be prepared.
____ 196. Bases taste
a. soapy.
b. sour.
c. sweet.
d. bitter.
____ 197. Bases feel
a. rough.
b. moist.
c. slippery.
d. dry.
____ 198. Bases react with
a. acids to produce salts and water.
b. salts to produce acids and water.
c. water to produce acids and salts.
d. neither acids, salts, nor water.
____ 199. Aqueous solutions of bases
a. contain only two different elements.
b. carry electricity.
c. have very high boiling points.
d. cannot be prepared.
____ 200. Which of the following is perchloric acid?
a. HClO
b. HClO2
c. HClO3
d. HClO4
____ 201. Which of the following is chlorous acid?
a. HClO
b. HClO2
c. HClO3
d. HClO4
____ 202. An acid ending with the suffix -ic produces an anion with the
a. suffix -ate.
c. prefix hydro-.
b. suffix -ite.
d. suffix -ous.
____ 203. Arrhenius theorized that an acid is a chemical compound that
a. increases the concentration of hydrogen ions when dissolved in water.
b. increases the concentration of hydroxide ions when dissolved in water.
c. decreases the concentration of hydrogen ions when dissolved in water.
d. decreases the concentration of hydroxide ions when dissolved in water.
____ 204. Arrhenius theorized that a base is a chemical compound that
a. increases the concentration of hydrogen ions when dissolved in water.
b. increases the concentration of hydroxide ions when dissolved in water.
c. decreases the concentration of hydrogen ions when dissolved in water.
d. decreases the concentration of hydroxide ions when dissolved in water.
____ 205. A substance that ionizes nearly completely in aqueous solutions and produces H3O+ is a
a. weak base.
c. weak acid.
b. strong base.
d. strong acid.
____ 206. The substances produced when KOH(aq) neutralizes HCl(aq) are
a. HClO(aq) and KH(aq).
c. H2O(l) and KCl(aq).
b. KH2O+(aq) and Cl–(aq).
d. H3O+(aq) and KCl(aq).
____ 207. Pure water contains
a. water molecules only.
b. hydronium ions only.
c. hydroxide ions only.
d. water molecules, hydronium ions, and hydroxide ions.
____ 208. Pure water partially breaks down into charged particles in a process called
a. hydration.
c. self-ionization.
b. hydrolysis.
d. dissociation.
____ 209. What is the product of H+ ion and OH– ion concentrations in water?
a. 10–28
c. 10–7
–14
b. 10
d. 55.4
____ 210. Which expression represents the pH of a solution?
a. log[H+]
c. log[OH–]
+
b. –log[H ]
d. –log[OH–]
____ 211. If [H+] of a solution is greater than [OH–], the solution
a. is always acidic.
c. is always neutral.
b. is always basic.
d. might be acidic, basic, or neutral.
____ 212. If [H+] of a solution is less than [OH–], the solution
a. is always acidic.
c. is always neutral.
b. is always basic.
d. might be acidic, basic, or neutral.
____ 213. The pH scale in general use ranges from
a. 0 to 1.
b. –1 to 1.
c. 0 to 7.
d. 0 to 14.
____ 214. The pH of an acidic solution is
a. less than 0.
b. less than 7.
c. greater than 7.
d. greater than 14.
____ 215. The pH of a basic solution is
a. less than 0.
b. less than 7.
c. greater than 7.
d. greater than 14.
____ 216. A water solution whose pH is 4
a. is always neutral.
b. is always basic.
c. is always acidic.
d. might be neutral, basic, or acidic.
____ 217. A water solution whose pH is 10
a. is always neutral.
b. is always basic.
c. is always acidic.
d. might be neutral, basic, or acidic.
____ 218. What is the pH of a 10–4 M HCl solution?
a. 4
b. 6
c. 8
d. 10
____ 219. What is the pH of a 10–5 M KOH solution?
a. 3
b. 5
c. 9
d. 11
____ 220. The pH of a solution is 9.0. What is its H+ concentration?
a. 1 10-9 M
c. 1 10-5 M
-7
b. 1 10 M
d. 9 M
____ 221. The pH of a solution is 10.00. What is its OH– concentration?
a. 1.0  10–10 M
c. 1.0  10–4 M
–7
b. 1.0  10 M
d. 10 M
____ 222. What process measures the amount of a solution of known concentration required to react with a measured
amount of a solution of unknown concentration?
a. autoprotolysis
c. neutralization
b. hydrolysis
d. titration
____ 223. An acid-base titration involves a
a. composition reaction.
b. neutralization reaction.
c. single-displacement reaction.
d. decomposition reaction.
____ 224. What is the molarity of an HCl solution if 50.0 mL is neutralized in a titration by 40.0 mL of 0.400 M NaOH?
a. 0.200 M
c. 0.320 M
b. 0.280 M
d. 0.500 M
____ 225. What is the molarity of an HCl solution if 125 mL is neutralized in a titration by 76.0 mL of 1.22 M KOH?
a. 0.371 M
c. 0.617 M
b. 0.455 M
d. 0.742 M