Download Chemistry Final Exam Review

Chemistry Final Exam Review
True/False
Indicate whether the sentence or statement is true or false.
____
1. A chemical formula indicates the relative number of atoms of each kind in a chemical compound.
____
2. For a monatomic cation the name ending gets changed to -ide.
____
3. Nomenclature is also described as a naming system.
____
4. The correct chemical formula for the compound containing aluminum and oxygen is Al 2O3.
____
5. A compound between a metal and a nonmetal uses the prefix naming system.
____
6. An ionic compound composed of a cation and an anion is referred to as a salt.
____
7. Oxidation number, oxidation charge and oxidation state all refer to the same property.
____
8. Transition metals can only have one oxidation state.
____
9. To convert from grams to mole you use the molar mass of the substance.
____ 10. To convert to molecules you use Avogadro's number.
____ 11. A binary compound is composed of only two elements.
____ 12. Avogadro's number is 2.066 x 1023.
____ 13. Empirical formulas are also called the simplest formula.
____ 14. Empirical formulas show the exact number of each element contained in a chemical compound.
____ 15. Anions are positively charged particles.
____ 16. A chemical reaction is the process by which one or more substances are changed into one or more different
substances.
____ 17. The substances on the right side of the arrow in a chemical equation are known as the reactants.
____ 18. The evolution of heat indicates that a chemical reaction has occurred.
____ 19. The formation of a precipitate is a physical change not a chemical change.
____ 20. The Law of Conservation of Mass must be satisfied in a chemical reaction.
____ 21. A chemical equation can be written in the form of a sentence.
____ 22. The sideways arrow means "yields."
____ 23. The triangle over the sideways arrow indicates that a catalyst was used in the reaction.
____ 24. A reversible chemical reaction means that the reaction can travel forwards or backwards.
____ 25. Subscripts are used to balance chemical reactions.
____ 26. A synthesis reaction contains two products.
____ 27. A decomposition reaction contains at least two products.
____ 28. A combustion reaction will always yield water as a product.
____ 29. A single displacement reaction has two compounds as the reactants.
____ 30. An acid-base reaction is really a double displacement reaction.
____ 31. Stoichiometry deals with the mass relationships in a chemical reaction.
____ 32. A mole-to-mole ratio converts between two substances in the same chemical reaction.
____ 33. A mole ratio is used to convert between moles and grams.
____ 34. A mole-to-mass calculation converts between the grams of one substance to the moles of another substance in
the chemical equation.
____ 35. A mass-to-mass calculation converts between the grams of two substances in a chemical equation.
____ 36. The molar mass is used in a mole-to-mole calculation.
____ 37. The theoretical yield is the amount produced in the laboratory.
____ 38. The kinetic molecular theory helps to explain the behavior of gases.
____ 39. Collisions between gas particles and between particles and container walls are elastic collisions.
____ 40. According to the KMT there is an attraction between particles.
____ 41. Gas particles have a high density and therefore a low compressibility.
____ 42. Particles colliding with walls of a container create pressure.
____ 43. There are 101.3 mm Hg in 1 atm.
____ 44. Boyle's law deals with volume and temperature.
____ 45. The combined gas law is composed of volume, pressure, and temperature.
____ 46. Solidification is another term for freezing.
____ 47. The solute is the substance that does the dissolving.
____ 48. A homogeneous mixture is also called a solution.
____ 49. An electrolyte will not conduct electricity.
____ 50. An Arrhenius acid produces hydrogen ions in solution.
____ 51. The pH of an acid is between 8 and 14.
____ 52. If a solution is neutral its pH is approximately 7.
____ 53. An acid will turn litmus blue.
Multiple Choice
Identify the letter of the choice that best completes the statement or answers the question.
____ 54. A chemical formula for a molecular compound represents the composition of
a. a molecule.
c. the ions that make up the compound.
b. an atom.
d. the crystal lattice.
____ 55. The formula for carbon dioxide, CO2, can represent
a. one molecule of carbon dioxide.
c. one molar mass of carbon dioxide.
b. 1 mol of carbon dioxide molecules.
d. all of the above.
____ 56. What is the formula for zinc fluoride?
a. ZnF
b. ZnF2
c. Zn2F
d. Zn2F 3
____ 57. What is the formula for the compound formed by calcium ions and chloride ions?
a. CaCl
c. CaCl3
b. Ca2Cl
d. CaCl2
____ 58. 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
____ 59. What is the formula for aluminum sulfate?
a. AlSO4
b. Al2SO4
c. Al2(SO4)3
d. Al(SO4)3
____ 60. What is the formula for tin(IV) chromate?
a. Sn(CrO4)4
b. Sn2(CrO4)2
c. Sn2(CrO4)4
d. Sn(CrO4)2
____ 61. Name the compound Ni(ClO3)2.
a. nickel chlorate
b. nickel chloride
c. nickel chlorite
d. nickel peroxide
____ 62. Name the compound Zn3(PO4)2.
a. zinc potassium oxide
b. trizinc polyoxide
c. zinc phosphate
d. zinc phosphite
____ 63. Name the compound CuCO3.
a. copper(I) carbonate
b. cupric trioxycarbide
c. cuprous carbide
d. copper(II) carbonate
____ 64. Name the compound N2O5.
a. dinickel pentoxide
b. dinitrogen pentoxide
c. neon oxide
d. nitric oxide
____ 65. What is the formula for silicon dioxide?
a. SO2
b. SiO2
c. Si2O
d. S2O
____ 66. What is the formula for diphosphorous pentoxide?
a. P2PeO5
c. P2O4
b. PO5
d. P2O5
____ 67. What is the formula for carbon disulfide?
a. CaS2
b. CS2
c. S2C
d. SC2
____ 68. What is the oxidation number of oxygen in most compounds?
a. –8
c. 0
b. –2
d. +1
____ 69. In a compound, the algebraic sum of the oxidation numbers of all atoms equals
a. 0.
c. 8.
b. 1.
d. the charge on the compound.
____ 70. What is the oxidation number of oxygen in H2O2?
a. –2
c. +2
b. –1
d. +4
____ 71. What is the oxidation number of hydrogen in H2O?
a. 0
c. +2
b. +1
d. +3
____ 72. What is the oxidation number of sulfur in H2SO4?
a. –2
c. +4
b. 0
d. +6
____ 73. What is the oxidation number of oxygen in CO2?
a. –4
c. 0
b. –2
d. +4
____ 74. Name the compound PBr5.
a. potassium hexabromide
b. phosphorus(V) pentabromide
c. phosphorus(V) bromide
d. phosphoric acid
____ 75. The molar mass of an element is the mass of one
a. atom of the element.
c. gram of the element.
b. liter of the element.
d. mole of the element.
____ 76. What is the formula mass of ethyl alcohol, C2H5OH?
a. 30.328 amu
c. 45.061 amu
b. 33.271 amu
d. 46.069 amu
____ 77. 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
____ 78. The molar mass of CS2 is 76.14 g/mol. How many grams of CS 2 are present in 10.00 mol?
a. 0.13 g
c. 10.00 g
b. 7.614 g
d. 761.4 g
____ 79. What is the percentage composition of CF 4?
a. 20% C, 80% F
b. 13.6% C, 86.4% F
c. 16.8% C, 83.2% F
d. 81% C, 19% F
____ 80. 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. 47.263% Cu, 52.737% Cl
____ 81. The percentage composition of sulfur in SO2 is about 50%. What is the percentage of oxygen in this
compound?
a. 25%
c. 75%
b. 50%
d. 90%
____ 82. What is the empirical formula for a compound that is 31.9% potassium, 28.9% chlorine, and 39.2% oxygen?
a. KClO2
c. K2Cl2O3
b. KClO3
d. K2Cl2O5
____ 83. A compound contains 64 g of O and 4 g of H. What is the empirical formula for this compound?
a. H2O
c. HO2
b. H2O2
d. HO
____ 84. The empirical formula and the formula mass of a compound are needed to determine the compound's
a. molecular formula.
c. lattice structure.
b. bond energy.
d. toxicity.
____ 85. 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
____ 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. Which observation does NOT indicate that a chemical reaction has occurred?
a. formation of a precipitate
c. evolution of heat and light
b. production of a gas
d. change in total mass of substances
____ 88. A solid produced by a chemical reaction in solution that separates from the solution is called
a. a precipitate.
c. a molecule.
b. a reactant.
d. the mass of the product.
____ 89. After the correct formula for a reactant in an equation has been written, the
a. subscripts are adjusted to balance the equation.
b. formula should not be changed.
c. same formula must appear as the product.
d. symbols in the formula must not appear on the product side of the equation.
____ 90. 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
____ 91. 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.
____ 92. 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.
____ 93. 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.
____ 94. 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.
b. C + O2  CO2
c. CO2  C + O2
d. 2C + O  CO2
____ 95. In an equation, the symbol for a substance in water solution is followed by
a. (1).
c. (aq).
b. (g).
d. (s).
____ 96. 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
____ 97. Which coefficients correctly balance the formula equation CaO + H 2O  Ca(OH)2?
a. 2, 1, 2
c. 1, 2, 1
b. 1, 2, 3
d. 1, 1, 1
____ 98. After the first steps in writing an equation, the equation is balanced by
a. adjusting subscripts to the formula(s).
b. adjusting coefficients to the smallest whole-number ratio.
c. changing the products formed.
d. making the number of reactants equal to the number of products.
____ 99. Which equation is NOT balanced?
a. 2H2 + O2  2H2O
b. 4H2 + 2O2  4H2O
c. H2 + H2 + O2  H2O + H2O
d. 2H2 + O2  H2O
____ 100. The reaction 2Mg(s) + O2(g)  2MgO(s) is a
a. synthesis reaction.
c. single-replacement reaction.
b. decomposition reaction.
d. double-replacement reaction.
____ 101. The reaction Mg(s) + 2HCl(aq)  H2(g) + MgCl2(aq) is a
a. composition reaction.
c. single-replacement reaction.
b. decomposition reaction.
d. double-replacement reaction.
____ 102. The reaction Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq) is a
a. double-replacement reaction.
c. decomposition reaction.
b. synthesis reaction.
d. combustion reaction.
____ 103. The reaction 2KClO3(s)  2KCl(s) + 3O2(g) is a(n)
a. synthesis reaction.
c. combustion reaction.
b. decomposition reaction.
d. ionic reaction.
____ 104. 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.
____ 105. The replacement of bromine by chlorine in a salt is an example of a single-replacement reaction by
a. halogens.
c. water.
b. sodium.
d. electrolysis.
____ 106. When a slightly soluble solid compound is produced in a double-replacement reaction, a
a. gas bubbles off.
c. combustion reaction takes place.
b. precipitate is formed.
d. halogen is produced.
____ 107. If chlorine gas is produced by halogen replacement, the other halogen in the reaction must be
a. bromine.
c. astatine.
b. iodine.
d. fluorine.
____ 108. 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.
____ 109. What is the balanced equation when aluminum reacts with copper(II) sulfate?
a. Al + Cu2S  Al2S + Cu
c. Al + CuSO4  AlSO4 + Cu
b. 2Al + 3CuSO4  Al2(SO4)3 + 3Cu
d. 2Al + Cu2SO4  Al2SO4 + 2Cu
____ 110. The ability of an element to react is the element's
a. valence.
c. stability.
b. activity.
d. electronegativity.
____ 111. 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.
____ 112. Magnesium bromide + chlorine yield
a. Mg and BrCl.
c. MgBrCl.
b. MgCl and Br2.
d. Mg(Cl)2 and Br2.
____ 113. A determination of the masses and number of moles of sulfur and oxygen in the compound sulfur dioxide
would be studied in
a. reaction stoichiometry.
c. chemical equilibrium.
b. chemical kinetics.
d. composition stoichiometry.
____ 114. 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.
____ 115. 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.
____ 116. Each of the four types of reaction stoichiometry problems requires using a
a. table of bond energies.
c. Lewis structure.
b. chart of electron configurations.
d. mole ratio.
____ 117. In the chemical reaction wA + xB  yC + zD, a comparison of the number of moles of A to the number of
moles of C would be a(n)
a. mass ratio.
c. electron ratio.
b. mole ratio.
d. energy proportion.
____ 118. In the reaction N2 + 3H2  2NH3, what is the mole ratio of nitrogen to ammonia?
a. 1:1
c. 1:3
b. 1:2
d. 2:3
____ 119. In the reaction Zn + H2SO4  ZnSO4 + H2, what is the mole ratio of zinc to sulfuric acid?
a. 1:6
c. 1:2
b. 1:1
d. 3:1
____ 120. The Haber process for producing ammonia commercially is represented by the equation N 2(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
____ 121. For the reaction 2H2 + O2  2H2O, how many grams of water are produced from 6.00 mol of hydrogen?
a. 2.00 g
c. 54.0 g
b. 6.00 g
d. 108 g
____ 122. For the reaction 2Fe + O2  2FeO, how many grams of iron oxide are produced from 8.00 mol of iron?
a. 71.8 g
c. 712 g
b. 574 g
d. 1310 g
____ 123. For the reaction HCl + NaOH  NaCl + H2O, how many moles of hydrochloric acid are required to produce
150. g of water?
a. 1.50 mol
c. 8.32 mol
b. 4.16 mol
d. 12.2 mol
____ 124. For the reaction 3Fe + 4H2O  Fe3O4 + 4H2, how many moles of iron oxide are produced from 500 g of iron?
a. 1 mol
c. 9 mol
b. 3 mol
d. 12 mol
____ 125. For the reaction 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. 44 g
d. 200. g
____ 126. For the reaction 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. 825 g
____ 127. 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.
____ 128. 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.
____ 129. To determine the limiting reactant in a chemical reaction, one must know the
a. available amount of one of the reactants. c. available amount of each reactant.
b. amount of product formed.
d. speed of the reaction.
____ 130. According to the kinetic-molecular theory, particles of matter are in motion in
a. gases only.
c. solids, liquids, and gases.
b. gases and liquids.
d. solids only.
____ 131. An ideal gas is an imaginary 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.
____ 132. Unlike in an ideal gas, in a real gas
a. all particles move in the same direction.
b. all particles have the same kinetic energy.
c. the particles cannot diffuse.
d. the particles exert attractive forces on each other.
____ 133. Which is an example of gas diffusion?
a. inflating a flat tire
b. the odor of perfume spreading throughout a room
c. a cylinder of oxygen stored under high pressure
d. All of the above
____ 134. What happens to the volume of a gas during compression?
a. The volume increases.
b. The volume decreases.
c. The volume remains constant.
d. It is impossible to tell because all gases are different.
____ 135. 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
____ 136. When does a real gas behave like an ideal gas?
a. when the particles are far apart
b. when the kinetic energy of the particles is low
c. when the pressure is high
d. when the gas is liquefied
____ 137. What does the constant bombardment of gas molecules against the inside walls of a container produce?
a. temperature
c. pressure
b. density
d. diffusion
____ 138. A pressure of 745 mm Hg equals
a. 745 torr.
b. 1 torr.
c. 1 pascal.
d. 745 pascal.
____ 139. Convert the pressure 0.75 atm to mm Hg.
a. 101.325 mm Hg
b. 430 mm Hg
c. 570 mm Hg
d. 760 mm Hg
____ 140. Standard temperature is exactly
a. 100ºC.
b. 273ºC.
c. 0ºC.
d. 0 K.
____ 141. Standard pressure is exactly
a. 1 atm.
b. 760 atm.
c. 101.325 atm.
d. 101 atm.
____ 142. 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.
____ 143. 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.
____ 144. If the temperature remains constant, V and P represent the original volume and pressure, and V' and P'
represent the new volume and pressure, what is the mathematical expression for Boyle's law?
a. P'V = V'P
c. V'P' = VP
b. VV' = PP'
d.
____ 145. Pressure and volume changes at a constant temperature can be calculated using
a. Boyle's law.
c. Kelvin's law.
b. Charles's law.
d. Dalton's law.
____ 146. 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
____ 147. 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
____ 148. If V is the original volume, V' is the new volume, T is the original Kelvin temperature, and T' is the new
Kelvin temperature, how is Charles's law expressed mathematically?
a.
c.
b.
d.
____ 149. 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
____ 150. If V, P, and T represent the original volume, pressure, and temperature in the correct units, and V', P', and T'
represent the new conditions, what is the combined gas law?
a.
c.
b.
d.
____ 151. Suppose that the pressure of 1.00 L of gas is 380. mm Hg when the temperature is 200. K. At what
temperature is the volume 2.00 L and the pressure 0.750 atm?
a. 1.00 K
c. 219ºC
b. 600. K
d. 67.0 K
____ 152. 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
____ 153. A mixture of four gases exerts a total pressure of 860 mm Hg. Gases A and B each exert 220 mm Hg. Gas C
exerts 110 mm Hg. What pressure is exerted by gas D?
a. 165 mm Hg
c. 860 mm Hg
b. 310 mm Hg
d. cannot be determined
____ 154. The standard molar volume of a gas at STP is
a. 22.4 L.
c. g-mol wt/22.4 L.
b. g/22.4 L.
d. 1 L.
____ 155. All of the following equations are statements of the ideal gas law except
a. P = nRTV
c.
b.
d.
____ 156. What is the value of the gas constant?
a.
0.0821
b. 0.0281 L·atm
c.
0.0281
d. 0.0821 mol·K
____ 157. Calculate the approximate volume of a 0.600 mol sample of gas at 15.0ºC and a pressure of 1.10 atm.
a. 12.9 L
c. 24.6 L
b. 22.4 L
d. 129 L
____ 158. 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
____ 159. A gas sample with a mass of 0.467 g is collected at 20.ºC and 732.5 mm Hg. The volume is 200. mL. What is
the molar mass of the gas?
a. 58 g/mol
c. 290 g/mol
b. 180 g/mol
d. 730 g/mol
____ 160. Compared with the particles in a gas, the particles in a liquid
a. have more energy.
c. move around less.
b. are larger.
d. are farther apart.
____ 161. What causes particles in a liquid to escape into a gas state?
a. high kinetic energy
c. surface tension
b. a freezing temperature
d. the combining of liquids
____ 162. What is the physical change of a liquid to a solid by the removal of heat?
a. solidification
c. freezing
b. particle arrangement
d. both a and c
____ 163. Which of the following is an amorphous solid?
a. ice
c. graphite
b. diamond
d. glass
____ 164. During boiling, the temperature of a liquid
a. remains constant.
b. increases.
c. decreases.
d. approaches the standard boiling point.
____ 165. What is the freezing point of water at standard pressure?
a. –10ºC
c. 4ºC
b. 0ºC
d. 32ºC
____ 166. What is the boiling point of water at standard pressure?
a. 100ºC
c. 212ºC
b. 112ºC
d. 200ºC
____ 167. To conduct electricity, a solution must contain
a. nonpolar molecules.
c. ions.
b. polar molecules.
d. free electrons.
____ 168. Increasing the surface area between solute and solvent
a. increases the rate of dissolution.
b. decreases the rate of dissolution.
c. has no effect on the rate of dissolution.
d. can increase, decrease, or have no effect on the rate of dissolution.
____ 169. Which of the following will dissolve most rapidly?
a. sugar cubes in cold water
c. powdered sugar in cold water
b. sugar cubes in hot water
d. powdered sugar in hot water
____ 170. If the amount of dissolved solute in a solution at a given temperature is greater than the amount that can
permanently remain in solution at that temperature, the solution is said to be
a. saturated.
c. supersaturated.
b. unsaturated.
d. diluted.
____ 171. What is the molarity of a solution that contains 125 g NaCl in 4.00 L solution?
a. 0.535 M
c. 8.56 M
b. 2.14 M
d. 31.3 M
____ 172. What is the molality of a solution that contains 31.0 g HCl in 5.00 kg water?
a. 0.062 m
c. 0.170 m
b. 0.425 m
d. 15.5 m
____ 173. Acids taste
a. sweet.
b. sour.
c. bitter.
d. salty.
____ 174. Acetic acid is found in significant quantities in
a. lemons.
c. sour milk.
b. vinegar.
d. apples.
____ 175. Acids make litmus paper turn
a. red.
b. yellow.
c. blue.
d. black.
____ 176. Bases taste
a. soapy.
b. sour.
c. sweet.
d. bitter.
____ 177. Bases feel
a. rough.
b. moist.
c. slippery.
d. dry.
____ 178. Which of the following is a binary acid?
a. H2SO4
b. CH3COOH
c. HBr
d. NaOH
____ 179. The name of a binary acid
a. has no prefix.
b. begins with the prefix bi-.
c. ends with the suffix -ous.
d. begins with the prefix hydro-.
____ 180. Which of the following is perchloric acid?
a. HClO
b. HClO2
c. HClO3
d. HClO4
____ 181. Which of the following is chlorous acid?
a. HClO
b. HClO2
c. HClO3
d. HClO4
____ 182. Which of the following is chloric acid?
a. HClO
b. HClO2
c. HClO3
d. HClO4
____ 183. Compared with acids that have the suffix -ic, acids that have the suffix -ous contain
a. more hydrogen.
c. less oxygen.
b. more oxygen.
d. the same amount of oxygen.
____ 184. According to the traditional definition, an acid contains
a. hydrogen and does not ionize.
b. hydrogen and ionizes to form hydrogen ions.
c. oxygen and ionizes to form hydroxide ions.
d. oxygen and ionizes to form oxygen ions.
____ 185. Strong acids are
a. strong electrolytes.
b. weak electrolytes.
c. nonelectrolytes.
d. nonionized.
____ 186. Which of the following is a strong acid?
a. HSO4–
b. H2SO4
c. CH3COOH
d. H3PO4
____ 187. Strong bases are
a. strong electrolytes.
b. weak electrolytes.
c. nonelectrolytes.
d. also strong acids.
____ 188. An amphoteric species is one that reacts as a(n)
a. acid only.
c. acid or base.
b. base only.
d. None of the above
____ 189. The pH of an acidic solution is
a. less than 0.
b. less than 7.
c. greater than 7.
d. greater than 14.
____ 190. The pH of a basic solution is
a. less than 0.
b. less than 7.
c. greater than 7.
d. greater than 14.
____ 191. What unknown quantity can be calculated after performing a titration?
a. volume
c. mass
b. concentration
d. density
____ 192. 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
____ 193. 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
Short Answer
Write the correct formula for the following compounds.
194. Aluminum iodide
195. Hydrofluoric acid
196. Phosphoric acid
197. Dinitrogen tetroxide
198. Carbon tetrachloride
199. Copper(II) sulfide
200. Chromium(VI) oxide
Write the correct name for the following compounds.
201. HBr
202. Cu2O
203. HNO2
204. N2O5
205. PbCl2
206. SO3
207. CO
208. NH4Br
209. K3PO3
Problem
Balance the following chemical reactions and identify the type.
210. Zn + HCl ZnCl2
211. Al4C3(s) + H2O(l)

CH4(g)
+
Al(OH)3(s)
212. Silver nitrate dissolved in water is added to potassium iodide also in water to form solid silver iodide and
aqueous potassium nitrate.
213. What mass in grams of sodium hydroxide is produced if 20.0 g of sodium metal reacts with excess water
according to the chemical equation 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)?
Chemistry Final Exam Review
Answer Section
TRUE/FALSE
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T
F
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F
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F
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F
F
T
F
T
F
T
T
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MULTIPLE CHOICE
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A
D
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D
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D
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B
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B
A
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A
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D
D
B
B
B
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B
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B
D
A
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A
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C
A
B
D
A
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B
D
B
B
B
A
C
A
C
C
A
C
C
D
A
C
D
A
C
D
B
D
B
A
A
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B
A
176.
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178.
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D
C
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D
D
B
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C
B
A
B
A
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B
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B
C
D
SHORT ANSWER
194. ANS:
AlI3
195. ANS:
HF
196. ANS:
H3PO4
197. ANS:
N2O4
198. ANS:
CCl4
199. ANS:
CuS
200. ANS:
CrO3
201. ANS:
Hydrobromic acid
202. ANS:
Copper(I) oxide
203. ANS:
Nitrous acid
204. ANS:
Dinitrogen pentoxide
205. ANS:
Lead(II) chloride
206. ANS:
Sulfur trioxide
207. ANS:
Carbon monoxide
208. ANS:
Ammonium bromide
209. ANS:
Potassium phosphite
PROBLEM
210. ANS:
1,2,1,1; Single Displacement
211. ANS:
1,12,3,4; double displacement
212. ANS:
AgNO3(aq) + KI(aq) --> AgI(s)
213. ANS:
34.8 g NaOH
+
KNO3(aq); DD