2015_Final Exam Study Guide

Multiple Choice Identify the letter of the choice that best completes the statement or answers the question.

CONTENT REVIEW

____ 1. What kind of energy is carried by objects in motion? a. kinetic energy c. radiant energy b. electrical energy d. potential energy

____ 2. Which temperature scales have units of equal size? a. Fahrenheit and Celsius c. Celsius and Kelvin b. Kelvin and Fahrenheit d. none of the above

____ 3. Which equation correctly relates the Celsius and Kelvin scales? a. K = °C – 273 c. K = °C + 100 b. °C = K + 273 d. °C = K – 273

____ 4. Which state of matter is characterized by definite shape and definite volume? a. plasma c. solid b. liquid d. gas

____ 5. Which of the following is not simply a change in state? a. melting c. freezing b. burning d. boiling

____ 6. Which of the following is a physical property? a. flammability c. tendency to rust b. density d. none of the above

____ 7. Which of the following is a physical change? a. cooking c. rusting b. burning d. crushing

____ 8. A blend of two or more pure substances is called a(n) a. plasma. c. mixture. b. compound. d. element.

____ 9. An atom is defined as the smallest part of an element that a. contains at least one proton, neutron, and electron. b. retains the chemical identity of that element. c. can carry an electric charge. d. is affected in a cathode ray tube.

____ 10. Which of the following statements is part of Dalton's atomic theory of matter? a. All atoms are identical. b. All atoms of a given element are identical. c. All atoms differ from one another. d. All atoms of a given element have the same mass.

____ 11. Approximately how many chemical elements are there? a. 30 c. 100 b. 70 d. 300

____ 12. A cathode ray consists of a. protons. c. neutrons. b. electrons. d. gamma rays.

____ 13. The experiment that revealed the charge of the electron involved the use of a. gold foil. c. oil droplets. b. photographic film. d. cathode ray tubes.

____ 14. What is the charge on an alpha particle? a. 2+ c. 1- b. 1+ d. 0

____ 15. The mass number of an atom is defined as its a. mass in amu. c. total number of protons and electrons. b. number of protons. d. total number of neutrons and protons.

____ 16. The symbol indicates a fluorine atom that contains

a. 19 protons and 9 neutrons c. 9 protons and 19 neutrons. b. 9 protons and 10 neutrons. d. 9 protons and 10 electrons.

____ 17. An alpha particle is the same as a(n) a. helium nucleus. c. electron. b. deuterium nucleus. d. neutron.

____ 18. Which wave property expresses how often the wave oscillates up and down? a. frequency c. wavelength b. amplitude d. speed

____ 19. The value 3.00 108 m/s expresses light's a. frequency. c. wavelength. b. amplitude. d. speed.

____ 20. Wavelength is defined as the distance between a. the trough and crest of a wave. b. the beginning and ends of two successive waves. c. the crest of one wave and the trough of another. d. successive crests of a wave.

____ 21. Which of the following has the longest wavelength? a. ultraviolet radiation c. X-rays b. infrared radiation d. gamma rays

____ 22. Which scientist explained how quantization could be used to account for the line spectra of

elements? a. de Broglie c. Einstein b. Bohr d. Heisenberg

____ 23. The quantum-mechanical model explains the atom by treating the a. atom as a point. b. electron as a particle. c. electron as a wave of quantized energy. d. electron as a wave of unquantized energy.

____ 24. How is an electron's principal quantum number symbolized? a. s c. n b. p d. d

____ 25. Under what conditions can two electrons occupy the same orbital? a. never b. if they have opposite spins c. if they have parallel spins d. if they have different principal quantum numbers

____ 26. Which of the following elements is an exception to the Aufbau principle? a. carbon c. iron b. copper d. sulfur

____ 27. The elements in Group 1A of the periodic table are called the a. halogens. c. alkaline earth metals. b. alkali metals. d. noble gases.

____ 28. The elements in Group 2A of the periodic table are called the a. halogens. c. alkaline earth metals. b. alkali metals. d. noble gases.

____ 29. What is an atom's ionization energy? a. the energy given off when an electron is gained b. the energy needed to gain an electron c. the energy given off when an electron is lost d. the energy needed to remove an electron

____ 30. When an atom becomes a negative ion, it a. loses protons. c. becomes larger. b. remains the same size. d. becomes smaller.

____ 31. Which of the following is not an alkali metal? a. strontium c. lithium b. francium d. rubidium

____ 32. Which of the following are found in the elemental state in nature? a. alkali metals only c. both alkali and alkaline earth metals b. alkaline earth metals only d. neither alkali nor alkaline earth metals

____ 33. An alloy is a a. highly pure metal. c. mixture of metals. b. highly pure nonmetal. d. mixture of nonmetals.

____ 34. What is the main component of steel? a. carbon c. chromium b. nickel d. iron

____ 35. An example of an actinide element is a. radium. c. thorium. b. iridium. d. cerium.

____ 36. Ionic compounds are normally a. good conductors of electricity in the solid state. b. gases at room temperature. c. formed when a metal transfers its valence electrons to a nonmetal. d. electrically charged.

____ 37. What is the general rule for writing the formula of a binary ionic compound? a. Use a ratio of cations to anions that gives equal numbers of each kind of ion.

b. Use the charge of each ion as its subscript in the formula. c. Use the charge of each ion as the subscript for the other ion. d. Use the simplest whole number ratio of cations and anions that will give an

electrically neutral compound.

____ 38. All of the following compounds are exceptions to the octet rule except a. BCl3. c. Br2. b. NO. d. SI4.

____ 39. A bond is classified as nonpolar covalent if the difference in the electronegativities between the 2

atoms is a. 2.1 or more. c. less than 0.4. b. between 0.5 and 2. d. less than zero.

____ 40. Which of the following does a structural formula reveal about chemical bonds? a. their arrangement in space c. both a and b b. their presence and type d. neither a nor b

____ 41. What shape does an ammonia molecule have? a. pyramidal c. tetrahedral b. trigonal planar d. bent

____ 42. In a molecule of carbon dioxide, what kinds of bonds connect the carbon atom to the 2 oxygen

atoms? a. two single bonds c. one double bond and one triple bond b. one single bond and one double bond d. two double bonds

____ 43. The shape of a methane molecule is a. pyramidal. c. tetrahedral. b. trigonal planar. d. linear.

____ 44. In a carbon dioxide molecule, the bond angle is a. 107°. c. 120°. b. 109.5°. d. 180°.

____ 45. How many unshared pairs of valence electrons does the central atom in a bent molecule have? a. none c. two b. one d. three

____ 46. How would you describe a molecule of carbon dioxide? a. nonpolar, with nonpolar bonds c. polar, with nonpolar bonds b. nonpolar, with polar bonds d. polar, with polar bonds

____ 47. A substance that is formed by a chemical reaction is called a a. mole. c. coefficient. b. product. d. reactant.

____ 48. In a chemical equation, the symbol that takes the place of the word yields is a(n) a. equal sign. c. plus sign. b. coefficient. d. arrow.

____ 49. In balancing a chemical equation, which of the following are you allowed to do? a. change subscripts c. change superscripts b. write coefficients d. add new substances

____ 50. Which of the following symbols means a substance is in water solution? a. (aq) c. (w)

b. (s) d. (l)

____ 51. In what kind of reaction do two or more reactants come together to form a single product? a. direct combination c. single replacement b. double replacement d. decomposition

____ 52. In what kind of reaction does a single compound break down into two or more smaller

compounds or elements? a. direct combination c. single replacement b. double replacement d. decomposition

____ 53. How many atoms are in a sample of an element whose mass in grams is numerically equal to the

atomic mass? a. 1 c. 1 1023 b. 6.02 d. 6.02 1023

____ 54. Avogadro's number a. equals 1. c. depends on the substance. b. equals 6.02 1023. d. depends on the number of moles.

____ 55. The percentage composition of water is a. 67% H, 33% O. c. 11% H, 89% O. b. 2 H, 1 O. d. 2 H, 16 O.

____ 56. The actual number of atoms of each element in a molecular compound is given by the a. formula mass. c. molecular formula. b. molar mass. d. empirical formula.

____ 57. What is the empirical formula of water? a. HO c. H2O2 b. H2O d. HO2

____ 58. In a balanced chemical equation, the coefficients represent the relative numbers of a. uncombined atoms. c. moles. b. molecules. d. all of these.

____ 59. What is the term that represents the sum of the atomic masses of the atoms in 1 mole of a

compound? a. atomic weight c. density b. relative mass d. molar mass

____ 60. What is the volume of a mole of a gas at standard temperature and pressure? a. 6.02 L c. 273 L b. 22.4 L d. varies with the type of gas

____ 61. When two reactants are available in the stoichiometric proportions represented by the chemical

equation, a. only the most massive reactant will be used up. b. neither of the reactants will be used up. c. both of the reactants will be used up. d. only the least massive of the reactants will be used up.

____ 62. The ability of gases to spread rapidly through other gases is explained by a. frequent inelastic collisions of the particles. b. the high density of gases. c. the high temperature of gases.

d. constant random motion of particles.

____ 63. Gases do not form liquids, unless the temperature is lowered, because of the a. extremely small size of the gas particles. b. relatively slow motion of the gas particles. c. weak attractive forces between particles. d. presence of repulsive forces between particles.

____ 64. The characteristics (variables) of gases needed to describe a gas completely include all of the

following except a. pressure. c. temperature. b. density. d. volume.

____ 65. The temperature scale that has absolute zero as its zero point is the a. Celsius scale. c. Kelvin scale. b. Rankin scale. d. Fahrenheit scale.

____ 66. The molar volume of a gas a. is always 22.4 liters. b. varies with the type of gas. c. is dependent upon the temperature and pressure of the gas. d. is directly proportional to the molar mass of the gas.

____ 67. The ideal gas law describes the behavior of real gases under a. all conditions of temperature and pressure. b. most ordinary conditions of temperature and pressure. c. low temperatures and high-pressure conditions. d. any specified temperature and pressure conditions.

____ 68. R, the gas constant in the ideal gas equation, has a numerical value of 0.0821. What are the

units? a. atm-liter/mol-K c. atm-m3/mol-K b. Pa-m3/mol-K d. Pa-liter/mol-K

____ 69. If the phases of matter are arranged in order of increasing disorder, the arrangement would be a. solid, liquid, gas. c. gas, liquid, solid. b. gas, solid, liquid. d. liquid, solid, gas.

____ 70. The state in which a substance exists at room temperature depends on the a. size of the particles. b. number of particles. c. strength of the attractive forces between particles. d. number of collisions between particles.

____ 71. The average kinetic energy of the particles in solids, liquids, and gases is most closely related to

the a. temperature. c. volume. b. density. d. pressure.

____ 72. Water has unique and unusual properties because of the a. linear shape of water molecules. b. small size of the water molecules. c. dispersion forces between water molecules. d. hydrogen bonds between polar water molecules.

____ 73. In solution formation, energy is needed to

a. separate the solute particles from each other. b. separate the solvent molecules from each other. c. warm the solvent and solute. d. Answers a and b are both correct.

____ 74. The energy involved in the formation of a solution will be a. given off if it is a water solution. b. absorbed if the solution is solid. c. exothermic if solvation forces are higher than the forces required to separate solute

particles and solvent molecules. d. Answers a and c are both correct.

____ 75. Which of the following is not an important factor influencing solubility? a. chemical nature of solute c. chemical nature of solvent b. temperature d. volume of solvent

____ 76. Ionic solids best dissolve liquid solvents that are a. polar. c. viscous. b. nonpolar. d. transparent.

____ 77. Which of the following's solubility is most affected by pressure? a. ionic solids c. gases b. supersaturated solutions d. alloys

____ 78. The molal freezing point constant, Kf, a. varies with temperature and pressure. b. has the same value as the boiling point constant. c. is independent of the solute used in making the solution. d. is a specific effect of a solute on a given solvent.

____ 79. Starting with a sample of pure nitrogen dioxide, when equilibrium is achieved between nitrogen

dioxide and dinitrogen tetroxide, what happens to the dark brown color of the nitrogen dioxide? a. It increases in intensity. c. It becomes colorless. b. It decreases in intensity. d. It becomes white.

____ 80. What symbol is used to indicate a reversible reaction? a. c. R b. d.

____ 81. What is the general relationship between the concentration and the rate of a forward reaction? a. Rate is greater when product concentration is greater. b. Rate is greater when reactant concentration is greater. c. Rate is greater when reactant concentration is lower. d. No relationship exists between concentration and rate.

____ 82. What symbol is used to denote concentration? a. ( ) c. b. [ ] d. C

____ 83. Which of the following are related by the law of mass action? a. forward and reverse rates and relative concentrations b. forward and reverse rates and mass of reactants and products c. relative concentrations and the equilibrium constant d. the equilibrium constant and forward and reverse rates

____ 84. An equilibrium position is a a. set of equilibrium concentrations. b. ratio of rates. c. description of the placement of substances. d. state in which the reaction has ceased to occur.

____ 85. When equilibrium lies to the left, Keq is a. roughly 1. c. much greater than 1. b. negative. d. very small.

____ 86. The symbol >> means a. greater than. c. proceeds to the right. b. much greater than. d. occurs at a high rate.

____ 87. The symbol that separates reactants and products in a solubility equilibrium expression is a a. +. c. . b. . d. .

____ 88. Which of the following is an equation that represents dissolution? a. XY (s) XY (l) c. XY (s) X+ (aq) + Y– (aq) b. XY (l) XY (s) d. XY (s) X+ (aq) + Y– (aq)

____ 89. The symbol (aq) is used to denote a solution that is a. saturated. c. ionic. b. unsaturated. d. aqueous.

____ 90. In a solubility equilibrium expression equal to Keq, where do the concentrations of reactants and

products appear? a. products in the numerator, reactants in the denominator b. reactants in the numerator, products in the denominator c. Only product concentrations appear. d. Only reactant concentrations appear.

____ 91. A precipitate will form if Q is a. equal to 1. c. less than Ksp. b. equal to Ksp. d. greater than Ksp.

____ 92. Which of the following is a molecular equation? a. A+ (aq) + B– (aq) + C+ (aq) + D– (aq) AD (s) + C+ (aq) + B– (aq) b. A+ (aq) + D– (aq) AD (s) c. AB (aq) + CD (aq) AD (s) + CB (aq) d. none of the above

____ 93. Which of the following is a net ionic equation? a. A+ (aq) + B– (aq) + C+ (aq) + D– (aq) AD (s) + C+ (aq) + B– (aq) b. A+ (aq) + D– (aq) AD (s) c. AB (aq) + CD (aq) AD (s) + CB (aq) d. none of the above

____ 94. When the concentration of an ion that is part of the solubility expression is added to a saturated

solution, what will occur? a. Nothing will occur. c. The solution will become unsaturated. b. The solubility will increase. d. Precipitation will occur.

____ 95. The ionic compound formed in an acid-base neutralization reaction is a(n) a. indicator. c. hydroxide.

b. hydride. d. salt.

____ 96. How would you describe the solution of the salt formed from the reaction of a weak acid and a

strong base? a. acidic c. neutral b. basic d. need more information

____ 97. The definition of pH is a. –log [H3O

+]. b. the percent of hydronium ions. c. the parts per million of hydronium ions. d. log [OH–].

____ 98. A buffer keeps a solution's pH constant by a. providing neutralizing ions. b. increasing the concentration of H2O molecules. c. changing the [H3O

+]. d. releasing or absorbing H+ ions.

____ 99. What determines a buffer's capacity to neutralize added base? a. its temperature c. the amount of conjugate base b. the amount of weak acid d. Answers b and c are correct.

____ 100. To determine the concentration of an unknown acid, one of the requirements is a. litmus paper. c. a standard base. b. universal pH paper. d. a standard acid.

____ 101. The indicator chosen for the titration of a weak acid with a strong base should change color at

a(n) a. acidic pH. c. neutral pH. b. basic pH. d. pH of zero.

____ 102. All nuclear decay is accompanied by emission of a. radiation. c. visible light. b. protons. d. electrons.

____ 103. Cosmic rays are streams of a. high-energy charged particles from outer space. b. high-energy charged particles from Earth that move into outer space. c. high-energy waves from outer space. d. low-energy charged particles from outer space.

____ 104. A radioisotope that is often used to date objects made from plant or animal remains is a. 87Rb. c. 12C. b. 16O. d. 14C.

____ 105. Which isotope is often used to study thyroid function? a. carbon-14 c. cobalt-60 b. uranium-235 d. iodine-131

____ 106. Which isotope is often used externally to destroy cancerous tissue? a. carbon-14 c. cobalt-60 b. uranium-235 d. iodine-131

____ 107. The splitting of a large nucleus into two smaller nuclei of approximately equal size is called a. nuclear fission. c. neutron capture.

b. nuclear fusion. d. radioactive decay.

____ 108. A continuous series of fission reactions is called a a. fusion reaction. c. thermonuclear reaction. b. decay series. d. chain reaction.

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 2-1

____ 109. In Figure 2-1, what form of energy will sunlight produce when it warms objects on Earth and

increases the random internal motion of their particles? a. chemical potential energy c. gravitational potential energy b. mechanical energy d. thermal energy

Figure 2-2

____ 110. How should the mercury level shown on the Kelvin thermometer in Figure 2-2 be labeled? a. 100 c. 273 b. 212 d. 373

Figure 2-3

____ 111. Which sample in Figure 2-3 shows a compound? a. sample 1 only c. sample 3 only b. sample 2 only d. none of the samples

____ 112. Which sample in Figure 2-3 shows an element? a. sample 1 only c. sample 3 only b. sample 2 only d. none of the samples

Figure 2-4

____ 113. Which of the processes shown in Figure 2-4 illustrate a physical change that is not also a change

of state? a. process 1 only c. process 3 only b. process 2 only d. processes 1 and 3

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 3-1

____ 114. What did Thomson measure in his experiments with a cathode ray tube similar to the one shown

in Figure 3-1? a. the degree to which the ray was deflected b. the speed of the ray c. the wavelength of the ray d. the frequency of the ray

Figure 3-2

____ 115. In Figure 3-2, which letter represents the path of beta radiation? a. a c. c b. b d. d

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements

Figure 4-1

____ 116. What is the wavelength of the wave shown in Figure 4-1? a. 300 nm c. 900 nm b. 600 nm d. 1200 nm

Figure 4-2

____ 117. In Figure 4-2, which letter corresponds to ultraviolet radiation? a. a c. c b. b d. d

____ 118. In Figure 4-2, which part of the spectrum has the greatest frequency? a. radio waves c. gamma rays

b. visible d. Each part has the same frequency.

Figure 4-3

____ 119. What is the maximum total number of electrons that could be held by the orbitals in Figure 4-3? a. 2 c. 8 b. 4 d. 14

Figure 4-4

____ 120. Which of the diagrams shown in Figure 4-4 violates the Pauli exclusion principle but not Hund's

rule or the Aufbau principle? a. a c. c b. b d. d

____ 121. Using Figure 4-4, determine the electron configuration for phosphorus (P, atomic number 15). a. 1s22s22p63p5 c. 1s22s22p63s5 b. 1s22s22p63s23p3 d. 1s22s22p63s13p4

____ 122. Using Figure 4-4, determine the orbital diagram for the p sublevel of the next element, sulfur (S,

atomic number 16). a.

c.

b.

d.

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Properties of X Group Elements

Element Atomic mass

(amu)

Density

(g/cm3)

Melting point

(ºC)

Boiling point

(ºC)

X 10 3 600

Y 4 200 800

Z 20 300

Figure 5-1

____ 123. Imagine that a new element, R, is discovered. It has an atomic mass of 5 amu, a density of 2

g/cm3, a melting point of 0°C, and a boiling point of 400°C. Where in the X group of elements in

Figure 5-1 does it belong? a. above element X c. below element Z b. between elements Y and Z d. It does not belong in the X group.

Sample Square From the Periodic Table

19

K

39.0983

[Ar] 4s1

Figure 5-2

____ 124. In Figure 5-2, [Ar]4s1 represents the a. full electron configuration. c. full mass designation. b. abbreviated electron configuration. d. block designation.

Figure 5-3

____ 125. Which numbered regions in Figure 5-3 contain the representative elements? a. 1 and 2 c. 1 and 3 b. 2 and 3 d. 1, 2, and 3

____ 126. Which lettered region in Figure 5-3 represents the semimetals? a. a c. c

b. b d. none of the regions

Figure 5-4

____ 127. In Figure 5-4, which period tends to have the highest ionization energy? The lowest? a. 1; 4 c. 2; 3 b. 4; 1 d. 3; 2

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Period

Melting point (ºC) Melting point (ºC)

Alkali metals Alkaline earth metals

2 180.5 1287

3 97.8 649

4 63.7 839

5 39.0 768

6 28.6 727

Figure 6-1

____ 128. Using Figure 6-1, what element has a melting point of 180.5°C? a. hydrogen c. beryllium b. cesium d. lithium

____ 129. Using Figure 6-1, what element has a melting point of 839°C? a. calcium c. magnesium b. potassium d. lithium

____ 130. Using Figure 6-1, predict the approximate melting point of the period-7 alkaline earth metal. a. 700°C c. 750°C b. 747°C d. 1300°C

____ 131. How do the melting points for the alkaline earth metals in Figure 6-1 compare overall to those of

the alkali metals? a. The alkaline earth melting points are lower. b. The melting points are approximately the same. c. The alkaline earth melting points are higher. d. There is no general pattern.

Figure 6-2

____ 132. Using Figure 6-2, estimate the approximate first ionization energy for francium (period 7). a. 250 c. 390 b. 350 d. 540

____ 133. If the second ionization energies of the alkali metals were plotted, how would they compare to

the first ionization energies plotted in Figure 6-2 for the same elements? a. The second ionization energies would be higher. b. The second ionization energies would be lower. c. The two sets of values would be essentially the same. d. No general pattern would exist.

Figure 6-3

____ 134. In Figure 6-3, which of the elements shown would be unlikely to form any compounds? a. Al c. P b. Si d. Ar

____ 135. Refer to Figure 6-3. What trend would you expect in the degree of metallic character of the

elements of period 3? a. Metallic character should decrease with atomic number. b. Metallic character should increase with atomic number. c. Metallic character should be greatest in the middle of the period. d. There should be no general trend.

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

1A 2A 3A 4A 5A 6A 7A 8A

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Br Kr

Rb Xe

Cs

Figure 7-1

____ 136. Use Figure 7-1 to determine the symbol of the ion that has the electron configuration

[Ar]3d104s24p6. a. K+ c. Br2– b. Rb+ d. Cl–

TABLE OF COMMON IONS

H1+ Al3+ NO31–

Ca2+ Fe3+ SO42–

Zn2+ Cl1– CO32–

Cu2+ OH1– PO43–

P3–

Figure 7-2

____ 137. Using Figure 7-2, determine the formula of the compound formed between calcium and

hydroxide ions. a. CaOH c. CaOH2 b. Ca(OH)2 d. Ca2OH

____ 138. Determine the formula of copper(II) chloride, using Figure 7-2. a. Cu2Cl c. CuCl b. (CuCl)2 d. CuCl2

____ 139. Determine the formula of iron(III) nitrate, using Figure 7-2. a. Fe(NO3)3 c. Fe3NO3 b. FeN d. None of these is correct.

Element Electronegativity

F 4.0

O 3.5

Cl 3.0

Br 2.8

C 2.5

H 2.1

P 2.1

B 2.0

Mg 1.2

Ca 1.0

Na 0.9

K 0.8

Figure 7-3

____ 140. Using the information in Figure 7-3, predict the physical state of the compound formed between

phosphorous and hydrogen, PH3, at the temperature of boiling water. a. probably a gas b. probably a liquid c. probably a solid d. impossible to determine from the given data

____ 141. Using Figure 7-3, which two elements would form the most ionic bond possible? a. F–F c. Ca–O b. F–O d. K–F

Figure 7-4

____ 142. In Figure 7-4, what type of bonding should replace letter a? a. covalent c. nonpolar covalent b. polar covalent d. ionic

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 8-1

____ 143. Which molecule in Figure 8-1 has a central atom with one unshared pair of electrons? a. molecule 1 c. molecule 3 b. molecule 2 d. molecule 5

____ 144. Which molecule in Figure 8-1 has a central atom with two unshared pairs of electrons? a. molecule 1 c. molecule 4 b. molecule 3 d. molecule 5

____ 145. Which of the molecules in Figure 8-1 has a central atom with no unshared pairs of electrons? a. molecules 1, 2, and 3 c. molecules 1, 4, and 5 b. molecules 2, 4, and 5 d. molecules 2, 3, and 5

Bond Length (nm)

Cl – Cl 0.198

I – I 0.266

C – C 0.154

C = C 0.134

Figure 8-2

____ 146. Refer to Figure 8-2. What is the approximate bond length of the astatine molecule, At2? a. 0.02 nm c. 0.23 nm b. 0.13 nm d. 0.30 nm

____ 147. Refer to Figure 8-2. How do the lengths of the Si–Si bond and C–C bond compare? a. The Si–Si bond is shorter. c. The two bonds are the same length. b. The Si–Si bond is longer. d. It is impossible to predict.

____ 148. What trend is suggested by the C–C and C=C bond lengths in Figure 8-2? a. Bond length decreases as the number of bonding electrons increases. b. Bond length increases as the number of bonding electrons increases. c. Bond length increases as the total number of atoms bonded to an atom increases. d. There is no trend regarding bond length.

Figure 8-3

____ 149. Toward what atom do the arrow-type symbols in Figure 8-3 always point? a. the more electronegative atom c. the atom of higher mass b. the less electronegative atom d. the atom of lower reactivity

____ 150. Which of the molecules in Figure 8-3 have polar bonds? a. X3Y only c. XZ2 and XYZ only b. XZ2 only d. all of them

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 9-1

____ 151. What formula equation describes the reaction in Figure 9-1? a. H2 + Br2 H2Br2 c. H2 + Br2 HBr b. 2 H + 2 Br 2 HBr d. 2 H + 2 Br H2Br2

____ 152. Identify the reactant(s) in Figure 9-1. a. hydrogen only c. hydrogen bromide only b. both hydrogen and bromine d. both bromine and hydrogen bromide

Figure 9-2

____ 153. What is the balanced equation for the reaction shown in Figure 9-2? a. 2 NO + O2 2 N2O c. NO + O2 2 NO2 b. 2 NO + O2 2 NO2 d. N2 + O2 2 NO2

Figure 9-3

____ 154. Which reaction in Figure 9-3 might represent a chemical change in which an atom joins with a

compound to produce a more complex compound? a. reaction a c. reaction c b. reaction b d. reaction d

Figure 9-4

____ 155. Which metals in Figure 9-4 would be displaced from solutions of their compounds if a strip of

tin metal were added? a. lead and copper only b. potassium, calcium, and magnesium only c. all the metals except tin d. none of the metals

____ 156. Which of the metals in Figure 9-4 should you add to a solution that contains compounds of both

calcium and lead if you wished to displace lead metal but not calcium metal? a. magnesium or tin b. copper c. potassium d. None of the listed metals would produce the desired effect.

____ 157. Which of the metals in Figure 9-4 should you add to a solution that contains compounds of both

potassium and magnesium if you wished to displace potassium metal but not magnesium metal? a. calcium b. copper only c. tin, lead, or copper d. None of the listed metals would produce the desired effect.

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 10-1

____ 158. Using Figure 10-1, determine how many atoms of carbon are present in 2 moles of ethanol. a. 2 c. 2.41 1024 b. 4 d. 6.02 1023

____ 159. Using Figure 10-1, what is the total mass of the carbon in 2 moles of butane? a. 48.0 g c. 48.0 amu

b. 96.0 g d. 96.0 amu

____ 160. What is the formula mass of the acetic acid molecule in Figure 10-1? a. 30.0 g c. 60.0 amu b. 30.0 amu d. 60.0 g

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 11-1

____ 161. Figure 11-1 shows the process of electrolysis. How many grams of hydrogen will be produced if

you begin with 1.62 g of water? a. 0.09 grams c. A gas cannot be measured in grams. b. 0.18 grams d. need more information

____ 162. Figure 11-1 illustrates the process of electrolysis. The two gases that collect at the electrodes are

combined in a container and the mixture is ignited. How will the volume of the water vapor

formed compare with the volumes of the two gases? a. The volume of water vapor will be equal to the volumes of the two gases

combined. b. The volume of the water vapor will be equal to the volume of the hydrogen gas

only. c. The volume of the water vapor will be equal to the volume of the oxygen gas only. d. The volume of the water vapor will be greater than the volume of the two gases

combined.

Figure 11-3

____ 163. The reaction of methane and oxygen is represented by the equation CH4 + 2 O2 2 H2O + CO2.

If the balloons of methane and oxygen in Figure 11-3 are forced into a third container,

thoroughly mixed, and then ignited, what substances will be present in the container after the

reaction is over? a. carbon dioxide and water vapor b. carbon dioxide, methane, oxygen, and water vapor c. carbon dioxide, water vapor, and oxygen d. carbon dioxide, water vapor, and methane

____ 164. The reaction of hydrogen and oxygen is represented by the equation 2 H2 + O22 H2O. If the

balloons of hydrogen and oxygen in Figure 11-3 are both forced into a third container,

thoroughly mixed, and ignited, what substances will be present in the container after the reaction

is over? a. hydrogen and water vapor c. hydrogen and oxygen b. oxygen and water vapor d. hydrogen, oxygen, and water vapor

Figure 11-4

____ 165. In Figure 11-4, what is the mole ratio for hydrogen reacting with chlorine? a. 1 + 1 c. 2:2 b. 2 + 2 d. 1:1

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 13-1

____ 166. At what altitude would a sample of gas in Figure 13-1 have a pressure of 100 kPa? a. sea level c. 50 km b. 1 km d. 100 km

Figure 13-3

____ 167. Using Figure 13-3, what would be the volume of a sample of gas at 500 K if its volume at 300 K

was 100 liters? a. 100 liters c. 167 liters b. 133 liters d. 200 liters

Figure 13-4

____ 168. In Figure 13-4, if the temperature is 0°C and the pressure is 1 atm, what volume of the gas

corresponds to Avogadro's number of particles? a. 1.0 liter c. 11.2 liters b. 6.0 liters d. 22.4 liters

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 14-1

Each dot in the graph above represents a hydrogen compound.

____ 169. In Figure 14-1, which compound is most likely to have the strongest intermolecular forces? a. a c. c b. b d. d

Figure 14-2

____ 170. How do the two temperatures represented by the curves T1 and T2 in Figure 14-2 compare? a. T1 is lower than T2 c. T1 equals T2 b. T1 is higher than T2 d. impossible to compare

Figure 14-3

____ 171. Using Figure 14-3, determine the approximate pressure exerted by the molecules of water vapor

at 50°C. a. 90 mm c. 20 mm b. 45 mm d. 0.0 mm

____ 172. If the atmospheric pressure in Figure 14-3 is lowered to 570 millimeters, what would be the

boiling temperature of water? a. 100°C c. 88°C b. 92°C d. 57°C

Figure 14-4

____ 173. In which region in Figure 14-4 will water exist as a liquid? a. a c. c b. b d. a, b, and c

____ 174. How could you convert a sample of water in Figure 14-4 at 10°C and 20 mm Hg to vapor

without decreasing the temperature?

a. Decrease the pressure to 8 mm Hg, so conditions for the sample fall on the solid

line. b. Decrease the pressure to 5 mm Hg, so conditions for the sample fall below the

solid line. c. Decrease the pressure to that of the triple point of water. d. It can't be done.

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 15-1

____ 175. Which substance in Figure 15-1 is the most soluble in water at 0 degrees C? a. NaNO3 c. NH3 b. KClO3 d. None are soluble at that temperature.

Kf and Kb Values for Several Solvents (ºC kg/mol)

solvent freezing point Kf boiling point Kb

water 0 1.86 100 0.51

benzene 5.5

5.1

80.1 2.53

camphor 178 40

208.2

5.95

Figure 15-3

____ 176. Use Figure 15-3 to find out how many moles of solute would be needed to lower the freezing

point of 1 kilogram of water by 4.65°C. a. 2.5 c. 4.65 b. 1.86 d. 8.6

____ 177. Which solvent in Figure 15-3 has its freezing and boiling points most affected by dissolved

solutes? a. water b. benzene c. camphor d. All would be affected an equal amount by equal concentrations.

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 16-1

____ 178. At what time is equilibrium reached in Figure 16-1? a. 0 minutes c. 2 minutes b. 1 minute d. 4 minutes

____ 179. How do the concentrations of X and Y in Figure 16-1 compare at equilibrium? a. The concentration of Y is greater than the concentration of X. b. The concentrations of X and Y are equal. c. The concentration of X is greater than the concentration of Y. d. X is not present at all at equilibrium.

____ 180. Refer to Figure 16-1. Does the equilibrium lie to the right or to the left? a. to the right c. neither to the right nor to the left b. to the left d. need more information

____ 181. If the value of Q were computed at time = 1 minute in Figure 16-1, how would it compare to the

value of Keq? a. Q would be greater than Keq. b. Q would be less than Keq. c. Q would be equal to Keq. d. The answer cannot be determined without more information.

Figure 16-2

____ 182. In Figure 16-2, does the equilibrium lie to the right or to the left? a. to the left b. to the right

c. neither to the right nor to the left d. The answer cannot be determined without more information.

____ 183. In Figure 16-2, how would the addition of A to the system at equilibrium affect the concentration

of AB, assuming the other conditions are kept constant? a. [AB] would increase. b. [AB] would decrease. c. [AB] would be unaffected. d. The answer cannot be determined without more information.

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Cation concentration(mol/L)

Substance in saturated solution Ksp

lead(II) sulfide, PbS 8.4 10–28

cadium hydroxide,

Cd(OH)2

1.4 10–5

silver carbonate, Ag2CO3 8.1 10–12

Figure 17-1

____ 184. What is the solubility equilibrium equation for lead(II) sulfide in Figure 17-1? a. PbS (s) Pb+ (aq) + S– (aq) b. PbS (s) Pb2+ (aq) + S2– (aq) c. 2 PbS (s) 2 Pb+ (aq) + S2– (aq) d. PbS (s) Pb2+ (aq) + 2 S2– (aq)

____ 185. What is the solubility product expression for silver carbonate in Figure 17-1? a. Ksp = [Ag+]2[CO3

2– ] c. Ksp = [Ag2+]2[CO32– ]

b. Ksp = [2 Ag+][CO32– ] d. Ksp = [Ag+][CO3

2– ]2

____ 186. In Figure 17-1, what is the value of Ksp for Cd(OH)2? a. 7.8 10–10 c. 1.1 10–14 b. 5.5 10–15 d. 2.7 10–15

____ 187. In Figure 17-1, what is the value of the silver ion concentration, in moles per liter, in saturated

solution? a. 2.5 10–4 c. 1.3 10–4 b. 4.0 10–6 d. 5.7 10–6

____ 188. Which of the three compounds in Figure 17-1 is the least soluble, in moles per liter? a. Ag2CO3 c. Cd(OH)2 b. PbS d. They are equally soluble.

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 18-1

____ 189. What particles will the solution in Figure 18-1 contain? a.

H+ ions and NO ions c. HNO3 formula units

b. HNO3 molecules d. HNO3 /H2O complexes

Acid Reaction Ka (AT 25°C)

hydrofluoric HF + H2O H3O+ + F– 6.6 10–4

formic HCOOH + H2O H3O + HCOO– 1.8 10–4

hydrocyanic HCN + H2O H3O + CN– 6.2 10–10

hydrogen carbonate ion HCO3– + H2O H3O

+ + CO32– 4.7 10–11

Figure 18-2

____ 190. Identify the strongest acid in Figure 18-2. a. HF c. HCO3

– b. F– d. CO3

2–

____ 191. Which acid in Figure 18-2 is least able to donate a proton? a. HF c. HCO3

– b. HCN d. CO3

2–

____ 192. Use Figure 18-2 to determine which 1 M solution, hydrofluoric acid or formic acid, would have

the greater concentration of hydronium ions. a. hydrofluoric acid c. neither b. formic acid d. need more information

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

[H3O+] pH

1.00 M 0

0.010 M 2

1.0 10–4 M 4

1.0 10–6 M 6

1.0 10–8 M 8

1.0 10–10 M 10

1.0 10–12 M 12

1.0 10–14 M 14

Figure 19-1

____ 193. Using Figure 19-1, determine the hydroxide ion concentration of a solution with a pH of 13. a. 1.0 10–13 c. 1

b. 13 d. 1.0 10–1

Figure 19-2

____ 194. What is the expected pH at point C in Figure 19-2? a. 4 c. 10 b. 7 d. impossible to estimate

____ 195. What are the relative strengths of the acid and base in the titration shown in Figure 19-2? a. strong acid, strong base c. strong acid, weak base b. weak acid, strong base d. weak acid, weak base

CONCEPT MASTERY

Use the diagrams to answer the questions or complete the statements.

Figure 24-1

____ 196. In Figure 24-1, what mass of the isotope is present at 0 minutes? a. 0 grams c. 12 grams

b. 6 grams d. 24 grams

____ 197. In Figure 24-1, what mass of the isotope is present at 40 minutes? a. 0 grams c. 4.5 grams b. 3 grams d. 6 grams

____ 198. How much of the isotope in Figure 24-1 remains at the end of two half-life periods? a. 48 grams c. 3 grams b. 6 grams d. 1.5 grams

____ 199. In Figure 24-1, what mass of the isotope will still be present at the end of four half-life periods? a. 0.75 grams c. 3 grams b. 1.5 grams d. 48 grams

Isotope

Type of

Decay

Isotope

produced

Half-life

H

beta

12.3 years

P

S

14.3 days

alpha

Po

3.8 days

Pu

alpha

2.4 104 years

Figure 24-2

____ 200. How many half-life periods are required for a 20-gram sample of Pu in Figure 24-2 to be

reduced to a mass of 5 grams? a. one c. three b. two d. four

Essay

PROBLEM SOLVING

Use the skills you have developed in this chapter to solve each problem.

201. Identify each of the underlined phrases in the following account of a laboratory procedure as

either a chemical change or a physical change. Explain your answers.

The unknown substance was tested to determine its properties. First, a sample was

ground up finely with a mortar and pestle. Next, samples were dissolved in water and

other solvents. A Bunsen burner was lighted and some of the pure chemical was placed

into a test tube and heated gently until it melted. Because the heating was too rapid, the

test tube cracked and the experiment was repeated. Heating was continued this time until

the substance boiled. Some fumes of the boiling substance reacted with oxygen in the air

to form a white powder

202. Lithium metal reacts with water to form aqueous lithium hydroxide, LiOH, and hydrogen gas.

Write a balanced chemical equation for the reaction, including abbreviations for the physical

states.

203. Determine the percentage composition of magnesium carbonate (MgCO3).

204. Ammonia gas (NH3) combines with oxygen gas (O2) to form nitrogen gas and water vapor. If 5.0

liters of ammonia are consumed in the reaction, how many liters of nitrogen are produced?

205. The pressure of a mixture of three gases in a fixed volume container is 300 kPa. If nitrogen

molecules are three times the number of hydrogen molecules and oxygen molecules are two

times the number of hydrogen molecules, what are the partial pressures of the three gases?

206. The concentration of Co2+ ions in a saturated solution of cobalt(II) hydroxide (Co(OH)2) is 4.0

10-6 M. Calculate the value of Ksp for this substance.

207. The Ka for HCOOH is 1.78 10–4. What would be the expected [H3O+] for a 1.0 M water

solution of HCOOH?

208. What is the hydroxide ion concentration in a 1.67 10–4 M HCl solution?

209. Choose a buffer for a solution that contains acetic acid, CH3COOH. Write an equation for the

addition of HCl and another for the addition of NaOH.

CRITICAL THINKING AND APPLICATION

Discuss each of the following in a brief paragraph.

210. What did Democritus contribute to the understanding of matter? Explain why his ideas were not

accepted at the time. Give one reason why they were reconsidered in the eighteenth century.

211. The gases in Group 8A of the periodic table are often referred to as the noble gases. Explain why

they were given that name and why the name may be inappropriate in terms of the properties of

those gases.

212. Explain why the formula of glucose is written as C6H12O6 and not as CH2O.

213. Explain why LiF (lithium fluoride) is a solid at room temperature and NO (nitric oxide) is a gas

at room temperature, even though their formula masses are quite close to each other (26 and 30,

respectively).

214. Why do the coefficients of gaseous substances in a chemical equation represent either numbers

of moles or numbers of liters?

215. How is the chemical reaction of burning wood in a wood stove controlled by limiting one of the

reactants?

216. Why is it useful for scientists to invent the idea of an ideal gas when no real gases are ideal in

their behavior?

217. Explain why evaporation of perspiration makes you feel cooler.

218. Describe how a solid can be a solution. Give two examples of a solid solution.

219. Explain why a human red blood cell shrivels when placed in a concentrated saltwater solution.

220. How is a reaction quotient used to obtain information about a reaction?

2015_Final Exam Study Guide

Answer Section

MULTIPLE CHOICE

1. A

2. C

3. D

4. C

5. B

6. B

7. D

8. C

9. B

10. B

11. C

12. B

13. C

14. A

15. D

16. B

17. A

18. A

19. D

20. D

21. B

22. B

23. C

24. C

25. B

26. B

27. B

28. C

29. D

30. C

31. A

32. D

33. C

34. D

35. C

36. C

37. D

38. C

39. C

40. B

41. A

42. D

43. C

44. D

45. C

46. B

47. B

48. D

49. B

50. A

51. A

52. D

53. D

54. B

55. C

56. C

57. B

58. D

59. D

60. B

61. C

62. D

63. C

64. B

65. C

66. C

67. B

68. A

69. A

70. C

71. A

72. D

73. D

74. C

75. D

76. A

77. C

78. D

79. B

80. A

81. B

82. B

83. C

84. A

85. D

86. B

87. A

88. C

89. D

90. A

91. D

92. C

93. B

94. D

95. D

96. B

97. A

98. D

99. B

100. C

101. B

102. A

103. A

104. D

105. D

106. C

107. A

108. D

109. D

110. D

111. A

112. C

113. A

114. A

115. B

116. B

117. C

118. C

119. C

120. B

121. B

122. A

123. A

124. B

125. C

126. B

127. A

128. D

129. A

130. A

131. C

132. B

133. A

134. D

135. A

136. B

137. B

138. D

139. A

140. A

141. D

142. D

143. C

144. A

145. B

146. D

147. B

148. A

149. A

150. D

151. C

152. B

153. B

154. C

155. A

156. A

157. D

158. C

159. B

160. C

161. B

162. B

163. D

164. B

165. D

166. B

167. C

168. D

169. A

170. A

171. A

172. C

173. B

174. B

175. C

176. A

177. C

178. D

179. C

180. A

181. B

182. A

183. D

184. A

185. A

186. C

187. A

188. D

189. A

190. A

191. D

192. A

193. D

194. B

195. A

196. C

197. B

198. C

199. A

200. B

ESSAY

201. The changes, in order, are physical (grinding does not change the identity of the substance),

physical (dissolving does not generally change the identity), chemical (burning involves changes

in chemical identity), physical (melting does not change the identity), physical (cracking does

not change the identity of the glass), physical (boiling does not change the substance's identity),

chemical (a chemical reaction is a change in identity).

202. 2 Li(s) + 2 H2O(l) 2 LiOH(aq) + H2(g)

203. Mg: 24.3 amu/atom 1 atom = 24.3 amu

C: 12.0 amu/atom 1 atom = 12.0 amu

O: 16.0 amu/atom 3 atoms = 48.0 amu

formula mass = 84.3 amu

molar mass = 84.3 g/mol

percentage of Mg: 100% = 28.8%

percentage of C: 100% = 14.2%

percentage of O: 100% = 56.9%

204. 2 NH3 + 3/2 O2 —> N2 + 3 H2O

Mole/liter ratio of N2 produced/NH3 used is 1:2.

1:2 = x:5, so x = 2.5 liters of N2 is produced.

205. Pressure is proportional to the number of particles.

Let x = number of hydrogen molecules

2x = number of oxygen molecules

3x = number of nitrogen molecules

6x = total number of molecules

partial pressure of nitrogen = 3/6 300 kPa = 150 kPa

partial pressure of oxygen = 2/6 300 kPa = 100 kPa

partial pressure of hydrogen = 1/6 300 kPa = 50 kPa

206. Co(OH)2 (s) Co2+ (aq) 2 OH– (aq)

[Co2+] = 4.0 10–6 M

[OH–] = 2[Co2+]= 2 4.0 10–6 M

= 8.0 10–6

Ksp = [Co2+][OH–]2 = 4.0 10–6 (8.0 10–6)2

= 2.6 10–16

207. 1.78 10-4 =

Assume X is small compared to 1.0.

1.78 10–4 = X2

X = 1.3 10–2 (Assumption introduces 1% error.)

208. [OH–] = 1 10–14/1.67 10–4

= 6.0 10–11 M

209. For the addition of HCl: CH3COO– + H3O+ CH3COOH + H2O

For the addition of NaOH: CH3COOH + OH– CH3COO– + H2O

210. Democritus proposed the idea that matter is made up of tiny indivisible particles called atoms.

His ideas were not accepted because he could not account for what held the atomic particles

together in matter. Later, the discovery of the laws of definite composition and constant

composition made chemists reexamine his ideas. (Either law is an acceptable answer.)

211. The noble gases received that name because they were thought to remain aloof, or apart, from

reaction with all other materials. However, this supposed inertness was discovered not to be the

case for several of the gases, which can react with elements, such as fluorine, to form

compounds.

212. The formula of glucose is written as C6H12O6 and not as CH2O because a molecule of glucose

contains 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms.

213. Lithium fluoride is a solid because the electronegativity difference between two elements is so

great that the bonding is ionic. Nitric oxide has a small enough electronegativity difference for

covalent bonding and thus will be a molecular substance with low molecular mass. LiF will exist

as an ionic solid, and NO will be a gas at room temperature.

214. The coefficients of the gaseous substances represent either the volume ratios or the mole ratios

because equal volumes of gases at the same temperature and pressure contain equal numbers of

particles.

215. The burning of wood is the chemical combination of substances in the wood with oxygen from

the air. By limiting the amount of air available, the reaction can either be stopped or slowed

down.

216. The ideal gas was invented to fit the ideal gas law and allows the behavior of real gases to be

approximated. Adjustments can be made to take into account the size of gas particles and the

presence of forces of attraction between the particles.

217. The particles that leave the body are the particles with the greatest kinetic energy. The resulting

average kinetic energy of the remaining particles is lowered, and this decrease is observed as a

decrease in temperature.

218. A solution is a homogenous mixture. Small amounts of a solid solute can be present in a solid

solvent. Sterling silver, gold jewelry, and other alloys are examples of solid solutions.

219. A red blood cell has a lower concentration of salt than the salt water and therefore osmotic

pressure will cause water to go from the red blood cell out to the salt water. As the red blood cell

loses water, it will shrivel and eventually rupture.

220. A reaction quotient, Q, is calculated by measuring the actual quantities of reactants and products

present in a system that may or may not be at equilibrium. The value of Q is then compared with

a known value of Keq for the system. If Q is less than Keq, the system has not yet reached

equilibrium, and more products will be formed. If Q is greater than Keq, the system contains too

much product to be at equilibrium, and a shift to the left will take place. If Q equals Keq, the

system is at equilibrium.