Bonding Practice Exam

Bonding Practice Exam Matching Match each item with the correct statement below. a. halide ion e. valence electron b. octet rule f. coordination number c. ionic bond g. metallic bond d. electron dot structure 1. an electron in the highest occupied energy level of an atom 2. Atoms react so as to acquire the stable electron structure of a noble gas. 3. a depiction of valence electrons around the symbol of an element 4. an anion of chlorine or other halogen 5. the force of attraction binding oppositely charged ions together 6. the attraction of valence electrons for metal ions 7. the number of ions of opposite charge surrounding each ion in a crystal Match each item with the correct statement below. a. coordinate covalent bond d. single covalent bond b. double covalent bond e. polar bond c. structural formula f. hydrogen bond 8. a depiction of the arrangement of atoms in molecules and polyatomic ions 9. a covalent bond in which only one pair of electrons is shared 10. a covalent bond in which two pairs of electrons are shared 11. a covalent bond in which the shared electron pair comes from only one of the atoms 12. a covalent bond between two atoms of significantly different electronegativities 13. a type of bond that is very important in determining the properties of water and of important biological molecules such as proteins and DNA Match each item with the correct statement below. a. network solid e. tetrahedral angle b. bonding orbital f. VSEPR theory c. dipole interaction g. sigma bond d. bond dissociation energy 14. energy needed to break a single bond between two covalently bonded atoms 15. symmetrical bond along the axis between the two nuclei 16. molecular orbital that can be occupied by two electrons of a covalent bond 17. 109.5 18. shapes adjust so valence-electron pairs are as far apart as possible 19. attraction between polar molecules 20. crystal in which all the atoms are covalently bonded to each other Multiple Choice Identify the letter of the choice that best completes the statement or answers the question. 21. How many valence electrons are in an atom of phosphorus? a. 2 c. 4 b. 3 d. 5 22. How many valence electrons are in an atom of magnesium? a. 2 c. 4

b. 3 d. 5 23. How many valence electrons does a helium atom have? a. 2 c. 4 b. 3 d. 5 24. How many valence electrons are in a silicon atom? a. 2 c. 6 b. 4 d. 8 25. What is the name given to the electrons in the highest occupied energy level of an atom? a. orbital electrons c. anions b. valence electrons d. cations 26. How does calcium obey the octet rule when reacting to form compounds? a. It gains electrons. b. It gives up electrons. c. It does not change its number of electrons. d. Calcium does not obey the octet rule. 27. What is the maximum charge an ion is likely to have? a. 2 c. 4 b. 3 d. 5 28. What is the electron configuration of the calcium ion? a. 1s 2s 2p 3s 3p c. 1s 2s 2p 3s 3p 4s b. 1s 2s 2p 3s 3p 4s d. 1s 2s 2p 3s 29. What is the electron configuration of the gallium ion? a. 1s 2s 2p 3s 3p c. 1s 2s 2p 3s 3p 4s 4p b. 1s 2s 2p 3s 3p 4s d. 1s 2s 2p 3s 3p 3d 30. What is the charge on the strontium ion? a. 2 c. 1 b. 1 d. 2 31. The octet rule states that, in chemical compounds, atoms tend to have. a. the electron configuration of a noble gas b. more protons than electrons c. eight electrons in their principal energy level d. more electrons than protons 32. How many electrons does silver have to give up in order to achieve a pseudo-noble-gas electron configuration? a. 1 c. 3 b. 2 d. 4 33. How many electrons does barium have to give up to achieve a noble-gas electron configuration? a. 1 c. 3 b. 2 d. 4 34. What is the formula of the ion formed when potassium achieves noble-gas electron configuration? a. K c. K

b. K d. K 35. Which of the following ions has a pseudo-noble-gas electron configuration? a. Fe c. Cu b. Mn d. Ni 36. Which of the following elements does NOT form an ion with a charge of 1? a. fluorine c. potassium b. hydrogen d. sodium 37. What is the formula of the ion formed when tin achieves a stable electron configuration? a. Sn c. Sn b. Sn d. Sn 38. What is the formula of the ion formed when cadmium achieves a pseudo-noble-gas electron configuration? a. Cd c. Cd b. Cd d. Cd 39. Which of the following is a pseudo-noble-gas electron configuration? a. 1s 2s 2p 3s 3d c. 1s 2s 2p 3s 3p 3d b. 1s 2s 2p 3s 3p d. 1s 2s 2p 3s 3d 4s 40. How many electrons does nitrogen gain in order to achieve a noble-gas electron configuration? a. 1 c. 3 b. 2 d. 4 41. What is the formula of the ion formed when phosphorus achieves a noble-gas electron configuration? a. P c. P b. P d. P 42. How does oxygen obey the octet rule when reacting to form compounds? a. It gains electrons. b. It gives up electrons. c. It does not change its number of electrons. d. Oxygen does not obey the octet rule. 43. The electron configuration of a fluoride ion, F, is. a. 1s 2s 2p c. 1s 2s 2p 3s b. the same as that of a neon atom d. the same as that of a potassium ion 44. What is the electron configuration of the oxide ion (O )? a. 1s 2s 2p c. 1s 2s b. 1s 2s 2p d. 1s 2s 2p 45. What is the electron configuration of the iodide ion? a. 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p b. 1s 2s 2p 3s 3p 3d 4s 4p 4d c. 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s

d. 1s 2s 2p 3s 3p 3d 4s 4p 46. What is the charge on the cation in the ionic compound sodium sulfide? a. 0 c. 2 b. 1 d. 3 47. Which of the following occurs in an ionic bond? a. Oppositely charged ions attract. b. Two atoms share two electrons. c. Two atoms share more than two electrons. d. Like-charged ions attract. 48. What is the net charge of the ionic compound calcium fluoride? a. 2 c. 0 b. 1 d. 1 49. A compound held together by ionic bonds is called a. a. diatomic molecule c. covalent molecule b. polar compound d. salt 50. Which of the following is true about an ionic compound? a. It is a salt. c. It is composed of anions and cations. b. It is held together by ionic bonds. d. all of the above 51. How many valence electrons are transferred from the nitrogen atom to potassium in the formation of the compound potassium nitride? a. 0 c. 2 b. 1 d. 3 52. How many valence electrons are transferred from the calcium atom to iodine in the formation of the compound calcium iodide? a. 0 c. 2 b. 1 d. 3 53. What is the formula unit of sodium nitride? a. NaN c. Na N b. Na N d. NaN 54. What is the formula unit of aluminum oxide? a. AlO c. AlO b. Al O d. Al O 55. What is the name of the ionic compound formed from lithium and bromine? a. lithium bromine c. lithium bromium b. lithium bromide d. lithium bromate 56. What is the formula for sodium sulfate? a. NaSO c. Na(SO ) b. Na SO d. Na (SO ) 57. What is the formula for potassium sulfide?

a. KS c. KS b. K S d. K S 58. Which of the following compounds has the formula KNO? a. potassium nitrate c. potassium nitrite b. potassium nitride d. potassium nitrogen oxide 59. Which of the following pairs of elements is most likely to form an ionic compound? a. magnesium and fluorine c. oxygen and chlorine b. nitrogen and sulfur d. sodium and aluminum 60. Ionic compounds are normally in which physical state at room temperature? a. solid c. gas b. liquid d. plasma 61. Which of the following is true about the melting temperature of potassium chloride? a. The melting temperature is relatively high. b. The melting temperature is variable and unpredictable. c. The melting temperature is relatively low. d. Potassium chloride does not melt. 62. What does the term coordination number in ionic crystals refer to? a. the total number of valence electrons in an atom b. the number of oppositely charged ions surrounding a particular ion c. the number of atoms in a particular formula unit d. the number of like-charged ions surrounding a particular ion 63. Under what conditions can potassium bromide conduct electricity? a. only when melted b. only when dissolved c. only when it is in crystal form d. only when melted or dissolved in water 64. Which of the following is NOT a characteristic of most ionic compounds? a. They are solids. b. They have low melting points. c. When melted, they conduct an electric current. d. They are composed of metallic and nonmetallic elements. 65. Which of the following particles are free to drift in metals? a. protons c. neutrons b. electrons d. cations 66. What is the basis of a metallic bond? a. the attraction of metal ions to mobile electrons b. the attraction between neutral metal atoms c. the neutralization of protons by electrons d. the attraction of oppositely charged ions 67. What characteristic of metals makes them good electrical conductors? a. They have mobile valence electrons. b. They have mobile protons.

c. They have mobile cations. d. Their crystal structures can be rearranged easily. 68. An ionic bond is a bond between. a. a cation and an anion c. the ions of two different metals b. valence electrons and cations d. the ions of two different nonmetals 69. In a hexagonal close-packed crystal, every atom (except those on the surface) has neighbors. a. 6 c. 12 b. 8 d. 10 70. Which metallic crystal structure has a coordination number of 8? a. body-centered cubic c. hexagonal close-packing b. face-centered cubic d. tetragonal 71. Which is a typical characteristic of an ionic compound? a. Electron pairs are shared among atoms. b. The ionic compound has a low solubility in water. c. The ionic compound is described as a molecule. d. The ionic compound has a high melting point. 72. What is shown by the structural formula of a molecule or polyatomic ion? a. the arrangement of bonded atoms c. the number of metallic bonds b. the number of ionic bonds d. the shapes of molecular orbitals 73. Which of these elements does not exist as a diatomic molecule? a. Ne c. H b. F d. I 74. How do atoms achieve noble-gas electron configurations in single covalent bonds? a. One atom completely loses two electrons to the other atom in the bond. b. Two atoms share two pairs of electrons. c. Two atoms share two electrons. d. Two atoms share one electron. 75. Why do atoms share electrons in covalent bonds? a. to become ions and attract each other b. to attain a noble-gas electron configuration c. to become more polar d. to increase their atomic numbers 76. Which of the following elements can form diatomic molecules held together by triple covalent bonds? a. carbon c. fluorine b. oxygen d. nitrogen 77. Which noble gas has the same electron configuration as the oxygen in a water molecule? a. helium c. argon b. neon d. xenon 78. Which elements can form diatomic molecules joined by a single covalent bond? a. hydrogen only b. halogens only c. halogens and members of the oxygen group only

d. hydrogen and the halogens only 79. Which of the following is the name given to the pairs of valence electrons that do not participate in bonding in diatomic oxygen molecules? a. unvalenced pair c. inner pair b. outer pair d. unshared pair 80. Which of the following electron configurations gives the correct arrangement of the four valence electrons of the carbon atom in the molecule methane (CH )? a. 2s 2p c. 2s 2p 3s b. 2s 2p 3s d. 2s 2p 81. Which of the following diatomic molecules is joined by a double covalent bond? a. c. b. d. 82. A molecule with a single covalent bond is. a. CO c. CO b. Cl d. N 83. When one atom contributes both bonding electrons in a single covalent bond, the bond is called a(n). a. one-sided covalent bond c. coordinate covalent bond b. unequal covalent bond d. ionic covalent bond 84. Once formed, how are coordinate covalent bonds different from other covalent bonds? a. They are stronger. c. They are weaker. b. They are more ionic in character. d. There is no difference. 85. When H forms a bond with H O to form the hydronium ion H O, this bond is called a coordinate covalent bond because. a. both bonding electrons come from the oxygen atom b. it forms an especially strong bond c. the electrons are equally shared d. the oxygen no longer has eight valence electrons 86. Which of the following bonds is the least reactive? a. C C c. O H b. H H d. H Cl 87. How many valid electron dot formulas having the same number of electron pairs for a molecule or ion can be written when a resonance structure occurs? a. 0 c. 2 only b. 1 only d. 2 or more 88. In which of the following compounds is the octet expanded to include 12 electrons? a. H S c. PCl b. PCl d. SF 89. How many electrons can occupy a single molecular orbital? a. 0 c. 2 b. 1 d. 4

90. Molecular orbital theory is based upon which of the following models of the atom? a. classical mechanical model c. quantum mechanical model b. Bohr model d. Democritus s model 91. A bond that is not symmetrical along the axis between two atomic nuclei is a(n). a. alpha bond c. pi bond b. sigma bond d. beta bond 92. How is a pair of molecular orbitals formed? a. by the splitting of a single atomic orbital b. by the reproduction of a single atomic orbital c. by the overlap of two atomic orbitals from the same atom d. by the overlap of two atomic orbitals from different atoms 93. The side-by-side overlap of p orbitals produces what kind of bond? a. alpha bond c. pi bond b. beta bond d. sigma bond 94. Where are the electrons most probably located in a molecular bonding orbital? a. anywhere in the orbital b. between the two atomic nuclei c. in stationary positions between the two atomic nuclei d. in circular orbits around each nucleus 95. Sigma bonds are formed as a result of the overlapping of which type(s) of atomic orbital(s)? a. s only c. d only b. p only d. s and p 96. Which of the following bond types is normally the weakest? a. sigma bond formed by the overlap of two s orbitals b. sigma bond formed by the overlap of two p orbitals c. sigma bond formed by the overlap of one s and one p orbital d. pi bond formed by the overlap of two p orbitals 97. According to VSEPR theory, molecules adjust their shapes to keep which of the following as far apart as possible? a. pairs of valence electrons c. mobile electrons b. inner shell electrons d. the electrons closest to the nuclei 98. The shape of the methane molecule is called. a. tetrahedral c. four-cornered b. square d. planar 99. What causes water molecules to have a bent shape, according to VSEPR theory? a. repulsive forces between unshared pairs of electrons b. interaction between the fixed orbitals of the unshared pairs of oxygen c. ionic attraction and repulsion d. the unusual location of the free electrons 100. Which of the following theories provides information concerning both molecular shape and molecular bonding? a. molecular orbital theory c. orbital hybridization theory

b. VSEPR theory d. Bohr atomic theory 101. Experimental evidence suggests that the H C H bond angles in ethene, C H, are. a. 90 c. 120 b. 109.5 d. 180 102. What type of hybrid orbital exists in the methane molecule? a. sp c. sp b. sp d. sp d 103. What is the shape of a molecule with a triple bond? a. tetrahedral c. bent b. pyramidal d. linear 104. What type of hybridization occurs in the orbitals of a carbon atom participating in a triple bond with another carbon atom? a. c. b. d. 105. How many pi bonds are formed when sp hybridization occurs in ethene, C H? a. 0 c. 2 b. 1 d. 3 106. Which of the following atoms acquires the most negative charge in a covalent bond with hydrogen? a. C c. O b. Na d. S 107. A bond formed between a silicon atom and an oxygen atom is likely to be. a. ionic c. polar covalent b. coordinate covalent d. nonpolar covalent 108. Which of the following covalent bonds is the most polar? a. H F c. H H b. H C d. H N 109. When placed between oppositely charged metal plates, the region of a water molecule attracted to the negative plate is the. a. hydrogen region of the molecule c. H O H plane of the molecule b. geometric center of the molecule d. oxygen region of the molecule 110. What is thought to cause the dispersion forces? a. attraction between ions c. sharing of electron pairs b. motion of electrons d. differences in electronegativity 111. Which of the forces of molecular attraction is the weakest? a. dipole interaction c. hydrogen bond b. dispersion d. single covalent bond 112. What causes dipole interactions? a. sharing of electron pairs b. attraction between polar molecules

c. bonding of a covalently bonded hydrogen to an unshared electron pair d. attraction between ions 113. What are the weakest attractions between molecules? a. ionic forces c. covalent forces b. Van der Waals forces d. hydrogen forces 114. What causes hydrogen bonding? a. attraction between ions b. motion of electrons c. sharing of electron pairs d. bonding of a covalently bonded hydrogen atom with an unshared electron pair 115. Why is hydrogen bonding only possible with hydrogen? a. Hydrogen s nucleus is electron deficient when it bonds with an electronegative atom. b. Hydrogen is the only atom that is the same size as an oxygen atom. c. Hydrogen is the most electronegative element. d. Hydrogen tends to form covalent bonds. 116. Which type of solid has the highest melting point? a. ionic solid c. metal b. network solid d. nonmetallic solid 117. What is required in order to melt a network solid? a. breaking Van der Waals bonds c. breaking hydrogen bonds b. breaking ionic bonds d. breaking covalent bonds Short Answer 118. Give the electron configurations for boron and its ion. 119. Give the electron configurations for mercury and its 2 ion. 120. Give the electron configuration for calcium the ion. 121. Give the electron configuration for the lithium ion. 122. Give the electron configurations for iodine and its 1 ion. 123. What is the formula for the oxide ion? 124. Give the electron configuration for the chloride ion. 125. Give the electron configuration for the oxide ion. 126. Write the formula for the compound barium oxide. 127. Write the formula for the compound rubidium phosphide. 128. Write the formula for the compound boron chloride. 129. Write the electron configuration diagram that shows the transfer of electrons that takes place to form the compound sodium fluoride. Include the electron configurations of the ions formed. Numeric Response 130. How many valence electrons are in rubidium? 131. How many valence electrons are in bromine? 132. What is the charge of a particle having 9 protons and 10 electrons? 133. How many electrons does a gallium atom give up when it becomes an ion? 134. What is the coordination number of both ions in the cesium chloride crystal? 135. How many valence electrons does an iodine atom have? 136. What is the total number of covalent bonds normally associated with a single carbon atom in a compound? 137. How many electrons are shared in a single covalent bond?

Essay 138. How many electrons does a nitrogen atom need to gain in order to attain a noble-gas electron configuration? 139. How many unshared pairs of electrons does the nitrogen atom in ammonia possess? 140. How many electrons does carbon need to gain in order to obtain a noble-gas electron configuration? 141. How many electrons are shared in a double covalent bond? 142. How many covalent bonds are in a covalently bonded molecule containing 1 phosphorus atom and 3 chlorine atoms? 143. How many unshared pairs of electrons are in a molecule of hydrogen iodide? 144. What is the bond angle in a water molecule? 145. Explain the octet rule and give an example of how it is used. 146. Explain what a pseudo-noble-gas electron configuration is. Give examples of ions that have this type of configuration. 147. Explain how atoms (ions) are held together in an ionic bond. Give an example of an ionic compound. 148. Why must each cation in an ionic solid be surrounded by anions? 149. Explain how a pure metal is held together. Include a definition of a metallic bond in your explanation. 150. Explain how scientists have used metallic bonding to account for many of the physical properties of metals, such as electrical conductivity and malleability. 151. What is bond dissociation energy, and how does it affect carbon compounds? 152. Can some atoms exceed the limits of the octet rule in bonding? If so, give an example. 153. Indicate how bonding is explained in terms of molecular orbitals. 154. Explain a pi bond and a sigma bond. Which of these bond types tends to be the weaker? Why? 155. Explain what is meant by VSEPR theory. Give an example of how VSEPR theory can be applied to predict the shape of a molecule. 156. Explain what is meant by orbital hybridization. Give an example of a molecule in which orbital hybridization occurs. 157. Explain what a polar molecule is. Provide an example. 158. What determines the degree of polarity in a bond? Distinguish between nonpolar covalent, polar covalent, and ionic bonds in terms of relative polarity. 159. What are dispersion forces? How is the strength of dispersion forces related to the number of electrons in a molecule? Give an example of molecules that are attracted to each other by dispersion forces. 160. Describe a network solid and give two examples.

Bonding Practice Exam Answer Section MATCHING 1. ANS: E DIF: L1 REF: p. 187 OBJ: 7.1.1 2. ANS: B DIF: L1 REF: p. 188 OBJ: 7.1.2 3. ANS: D DIF: L1 REF: p. 188 OBJ: 7.1.4 4. ANS: A DIF: L1 REF: p. 192 OBJ: 7.1.4 5. ANS: C DIF: L1 REF: p. 194 OBJ: 7.2.1 6. ANS: G DIF: L1 REF: p. 201 OBJ: 7.2.1 7. ANS: F DIF: L1 REF: p. 198 OBJ: 7.2.2 8. ANS: C DIF: L1 REF: p. 218 OBJ: 8.1.2 9. ANS: D DIF: L1 REF: p. 217 OBJ: 8.2.1 10. ANS: B DIF: L1 REF: p. 221 OBJ: 8.2.1 11. ANS: A DIF: L1 REF: p. 223 OBJ: 8.2.4 12. ANS: E DIF: L1 REF: p. 238 OBJ: 8.4.1 13. ANS: F DIF: L1 REF: p. 241 OBJ: 8.4.3 14. ANS: D DIF: L1 REF: p. 226 OBJ: 8.2.5 15. ANS: G DIF: L1 REF: p. 230 OBJ: 8.3.1 16. ANS: B DIF: L1 REF: p. 230 OBJ: 8.3.1 17. ANS: E DIF: L1 REF: p. 232 OBJ: 8.3.2 18. ANS: F DIF: L1 REF: p. 232 OBJ: 8.3.2 19. ANS: C DIF: L1 REF: p. 240 OBJ: 8.4.3 20. ANS: A DIF: L1 REF: p. 243 OBJ: 8.4.4 MULTIPLE CHOICE 21. ANS: D DIF: L1 REF: p. 187 OBJ: 7.1.1 22. ANS: A DIF: L1 REF: p. 188 OBJ: 7.1.1 23. ANS: A DIF: L1 REF: p. 188 OBJ: 7.1.1 24. ANS: B DIF: L1 REF: p. 188 OBJ: 7.1.1 25. ANS: B DIF: L1 REF: p. 187 OBJ: 7.1.1 26. ANS: B DIF: L1 REF: p. 188 OBJ: 7.1.1 27. ANS: B DIF: L1 REF: p. 189 OBJ: 7.1.1 28. ANS: A DIF: L2 REF: p. 188, p. 189 29. ANS: D DIF: L2 REF: p. 190 OBJ: 7.1.1 30. ANS: D DIF: L1 REF: p. 190 OBJ: 7.1.2 31. ANS: A DIF: L2 REF: p. 188 OBJ: 7.1.2 32. ANS: A DIF: L1 REF: p. 190 OBJ: 7.1.3 33. ANS: B DIF: L1 REF: p. 190 OBJ: 7.1.3 34. ANS: B DIF: L1 REF: p. 190 OBJ: 7.1.3 35. ANS: C DIF: L1 REF: p. 190 OBJ: 7.1.3 36. ANS: A DIF: L1 REF: p. 190 OBJ: 7.1.3 37. ANS: A DIF: L2 REF: p. 190 OBJ: 7.1.3 38. ANS: B DIF: L2 REF: p. 190 OBJ: 7.1.3 39. ANS: C DIF: L2 REF: p. 190 OBJ: 7.1.3 40. ANS: C DIF: L1 REF: p. 192 OBJ: 7.1.4 41. ANS: D DIF: L1 REF: p. 192 OBJ: 7.1.4 42. ANS: A DIF: L1 REF: p. 191 OBJ: 7.1.4 43. ANS: B DIF: L1 REF: p. 192 OBJ: 7.1.4 44. ANS: B DIF: L2 REF: p. 192 OBJ: 7.1.4 45. ANS: A DIF: L2 REF: p. 192 OBJ: 7.1.4 46. ANS: B DIF: L1 REF: p. 194 OBJ: 7.2.1 47. ANS: A DIF: L1 REF: p. 194 OBJ: 7.2.1 48. ANS: C DIF: L1 REF: p. 194 OBJ: 7.2.1 49. ANS: D DIF: L1 REF: p. 194 OBJ: 7.2.1 50. ANS: D DIF: L1 REF: p. 194 OBJ: 7.2.1 51. ANS: A DIF: L2 REF: p. 194 OBJ: 7.2.1 52. ANS: C DIF: L2 REF: p. 194 OBJ: 7.2.1 53. ANS: C DIF: L2 REF: p. 195 OBJ: 7.2.1 54. ANS: D DIF: L2 REF: p. 195 OBJ: 7.2.1 55. ANS: B DIF: L2 REF: p. 192, p. 195 56. ANS: B DIF: L2 REF: p. 192, p. 195 57. ANS: B DIF: L2 REF: p. 192, p. 195 58. ANS: A DIF: L2 REF: p. 192, p. 194 59. ANS: A DIF: L3 REF: p. 194 OBJ: 7.2.1

60. ANS: A DIF: L1 REF: p. 196 OBJ: 7.2.2 61. ANS: A DIF: L1 REF: p. 196 OBJ: 7.2.2 62. ANS: B DIF: L1 REF: p. 198 OBJ: 7.2.2 63. ANS: D DIF: L1 REF: p. 198 OBJ: 7.2.2 64. ANS: B DIF: L1 REF: p. 196, p. 198 65. ANS: B DIF: L1 REF: p. 201 OBJ: 7.3.1 66. ANS: A DIF: L1 REF: p. 201 OBJ: 7.3.1 67. ANS: A DIF: L1 REF: p. 201 OBJ: 7.3.1 68. ANS: A DIF: L1 REF: p. 201 OBJ: 7.2.1, 7.3.1 69. ANS: C DIF: L1 REF: p. 202 OBJ: 7.3.2 70. ANS: A DIF: L1 REF: p. 202 OBJ: 7.3.2 71. ANS: D DIF: L2 REF: p. 244 OBJ: 8.1.1 72. ANS: A DIF: L1 REF: p. 215 OBJ: 8.1.2 73. ANS: A DIF: L1 REF: p. 217 OBJ: 8.2.1 74. ANS: C DIF: L2 REF: p. 217 OBJ: 8.2.1 75. ANS: B DIF: L2 REF: p. 217 OBJ: 8.2.1 76. ANS: D DIF: L2 REF: p. 221 OBJ: 8.2.1 77. ANS: B DIF: L2 REF: p. 218 OBJ: 8.2.1 78. ANS: D DIF: L3 REF: p. 217, p. 218 79. ANS: D DIF: L1 REF: p. 218 OBJ: 8.2.2 80. ANS: D DIF: L3 REF: p. 220, p. 234 81. ANS: A DIF: L2 REF: p. 221 OBJ: 8.2.3 82. ANS: B DIF: L2 REF: p. 222 OBJ: 8.2.1, 8.2.4 83. ANS: C DIF: L2 REF: p. 223 OBJ: 8.2.4 84. ANS: D DIF: L2 REF: p. 223 OBJ: 8.2.4 85. ANS: A DIF: L2 REF: p. 225 OBJ: 8.2.4 86. ANS: B DIF: L2 REF: p. 226 OBJ: 8.2.5 87. ANS: A DIF: L1 REF: p. 227 OBJ: 8.2.6 88. ANS: D DIF: L2 REF: p. 229 OBJ: 8.2.7 89. ANS: C DIF: L1 REF: p. 230 OBJ: 8.3.1 90. ANS: C DIF: L1 REF: p. 230 OBJ: 8.3.1 91. ANS: C DIF: L1 REF: p. 230 OBJ: 8.3.1 92. ANS: D DIF: L2 REF: p. 230 OBJ: 8.3.1 93. ANS: C DIF: L2 REF: p. 231 OBJ: 8.3.1 94. ANS: B DIF: L2 REF: p. 231 OBJ: 8.3.1 95. ANS: D DIF: L2 REF: p. 230, p. 231 96. ANS: D DIF: L2 REF: p. 231 OBJ: 8.3.1 97. ANS: A DIF: L1 REF: p. 232 OBJ: 8.3.2 98. ANS: A DIF: L1 REF: p. 232 OBJ: 8.3.2 99. ANS: A DIF: L2 REF: p. 233 OBJ: 8.3.2 100. ANS: C DIF: L1 REF: p. 234 OBJ: 8.3.3 101. ANS: C DIF: L1 REF: p. 235 OBJ: 8.3.3 102. ANS: C DIF: L2 REF: p. 234 OBJ: 8.3.3 103. ANS: D DIF: L2 REF: p. 235 OBJ: 8.3.3 104. ANS: C DIF: L2 REF: p. 235 OBJ: 8.3.3 105. ANS: B DIF: L2 REF: p. 235 OBJ: 8.3.3 106. ANS: C DIF: L2 REF: p. 238, p. 239 107. ANS: C DIF: L2 REF: p. 238, p. 239 108. ANS: A DIF: L3 REF: p. 238, p. 239 109. ANS: A DIF: L3 REF: p. 239 OBJ: 8.4.2 110. ANS: B DIF: L1 REF: p. 240 OBJ: 8.4.3 111. ANS: B DIF: L1 REF: p. 240 OBJ: 8.4.3 112. ANS: B DIF: L1 REF: p. 240 OBJ: 8.1.1, 8.4.3 113. ANS: B DIF: L1 REF: p. 240 OBJ: 8.4.3 114. ANS: D DIF: L2 REF: p. 241 OBJ: 8.4.3 115. ANS: A DIF: L2 REF: p. 241 OBJ: 8.4.1, 8.4.3 116. ANS: B DIF: L1 REF: p. 243 OBJ: 8.4.4 117. ANS: D DIF: L1 REF: p. 243 OBJ: 8.4.4 SHORT ANSWER 118. B: 1s 2s 2p B : 1s DIF: L2 REF: p. 190 OBJ: 7.1.3

119. Hg: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s Hg : 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d DIF: L2 REF: p. 190 OBJ: 7.1.3 120. 1s 2s 2p 3s 3p DIF: L2 REF: p. 190 OBJ: 7.1.3 121. 1s DIF: L2 REF: p. 190 OBJ: 7.1.3 122. I: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p I : 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p DIF: L2 REF: p. 192 OBJ: 7.1.4 123. O DIF: L2 REF: p. 192 OBJ: 7.1.4 124. 1s 2s 2p 3s 3p DIF: L2 REF: p. 192 OBJ: 7.1.4 125. 1s 2s 2p DIF: L2 REF: p. 192 OBJ: 7.1.4 126. BaO DIF: L2 REF: p. 195 OBJ: 7.2.1 127. Rb P DIF: L2 REF: p. 195 OBJ: 7.2.1 128. BCl DIF: L2 REF: p. 195 OBJ: 7.2.1 129. Na(1s 2s 2p 3s ) F(1s 2s 2p ) Na (1s 2s 2p ) F (1s 2s 2p ) DIF: L3 REF: p. 195 OBJ: 7.2.1 NUMERIC RESPONSE 130. 1 DIF: L1 REF: p. 188 OBJ: 7.1.1 131. 7 DIF: L1 REF: p. 188 OBJ: 7.1.1 132. 1 DIF: L1 REF: p. 190 OBJ: 7.1.1 133. 3 DIF: L2 REF: p. 190 OBJ: 7.1.3 134. 8 DIF: L2 REF: p. 198 OBJ: 7.2.2 135. 7 DIF: L1 REF: p. 218 OBJ: 8.2.1 136. 4 DIF: L1 REF: p. 219 OBJ: 8.2.1 137. 2 DIF: L2 REF: p. 217 OBJ: 8.2.1 138. 3 DIF: L2 REF: p. 219 OBJ: 8.2.1 139. 1 DIF: L2 REF: p. 219 OBJ: 8.2.2 140. 4 DIF: L2 REF: p. 219 OBJ: 8.2.2 141. 4 DIF: L2 REF: p. 221 OBJ: 8.2.2 142. 3 DIF: L3 REF: p. 219 OBJ: 8.2.2 143. 3 DIF: L3 REF: p. 220 OBJ: 8.2.2 144. 105 DIF: L2 REF: p. 233 OBJ: 8.3.2 ESSAY 145. electron configurations of noble gases extremely stable. Octet rule in chemical reactions, elements gain or lose electrons to achieve a noble gas configuration. Called an octet because it consists of 8 valence electrons (s p ), 2 from the outermost s orbital and 6 from the outermost p orbital. Oxygen has 1s 2s 2p. When oxygen reacts to form ionic compounds, it completes its octet by gaining two electrons from the element it reacts with. These two electrons add to the p orbital of oxygen, giving it the electron configuration (1s 2p ) of neon. DIF: L3 REF: p. 188 OBJ: 7.1.2 146. A pseudo-noble-gas configuration has the form s p d. It has 18 electrons in the outer energy level and is a relatively stable configuration. Examples of ions with this configuration are Ag, Cu, Cd, and Hg. DIF: L3 REF: p. 190 OBJ: 7.1.3 2s

147. In an ionic bond, oppositely charged ions held together by electronic force of attraction between oppositely charged particles. Anions and cations present in a ratio that causes the total charge on the compound to be zero. Sodium phosphide, Na P, has three sodium ions for each phosphide ion. This ratio insures a zero total charge given the charges on the two individual ions (Na = 1, P = 3 ). DIF: L3 REF: p. 194 OBJ: 7.2.1 148. Like-charged ions are shielded from each other and electronic repulsion is reduced. Also, the force of attraction between oppositely charged ions is maximized. Each of these contributes to a lowering of energy and an increase in stability of the ionic compound. DIF: L3 REF: p. 196 OBJ: 7.2.2 149. Pure metal, like copper or iron, consists not of metal atoms, but of closely packed cations that are surrounded by mobile valence electrons that are free to drift from one part of the metal to another. Metallic bonds result from the attraction between the free-floating valence electrons and the positively charged metal ions. DIF: L3 REF: p. 201 OBJ: 7.3.1 150. Metallic bonds are the result of the attraction of free-floating valence electrons to positively charged metal ions. An electric current is a flow of electrons. As electrons enter one end of a piece of metal, some of the free-floating electrons leave the other end. Thus metals are good conductors of electricity. The cations in a piece of metal are insulated from each other by the free electrons. Thus when the metal is struck, the cations slide past each other easily. This makes the metal malleable and ductile. DIF: L3 REF: p. 201, p. 202 OBJ: 7.3.1 151. Bond dissociation energy is the energy required to break a single bond. The greater the bond dissociation energy, the more stable the compound. Due in part to the high bond dissociation energy of carbon-carbon bonds, carbon compounds are not very reactive chemically. DIF: L2 REF: p. 226 OBJ: 8.2.5 152. Yes, sulfur and phosphorus can expand the octet. They can have 12 or 10 valence electrons, respectively, when combined with small halogens. In PCl, phosphorus has 10 valence electrons. DIF: L2 REF: p. 228, p. 229 OBJ: 8.2.7 153. When two atoms combine, the overlap of their atomic orbitals produces molecular orbitals. An atomic orbital belongs to a particular atom, whereas a molecular orbital belongs to a molecule as a whole. Much like an atomic orbital, two electrons are required to fill a molecular orbital. A bonding orbital is a molecular orbital occupied by the two electrons of a covalent bond. DIF: L3 REF: p. 230 OBJ: 8.3.1 154. A pi bond is the bond formed as a result of the side-by-side overlap of two p orbitals. A sigma bond is the bond that results from a combination of two s orbitals, two p orbitals, or a p and an s orbital. Orbital overlap in pi bonding is less extensive than that for sigma bonding. Therefore, pi bonds tend to be weaker than sigma bonds. DIF: L3 REF: p. 230, p. 231 OBJ: 8.3.1 155. VSEPR (valence-shell electron-pair repulsion) theory states that because electron pairs repel, molecules adjust their shapes so that the valence-electron pairs, both bonding and non-bonding, are as far apart as possible. Methane, CH, for example, has four bonding electron pairs and no unshared pairs. The bonding pairs are farthest apart when the angle between the central carbon and each of its attached hydrogens is 109.5. This is the angle that is observed experimentally. DIF: L3 REF: p. 232 OBJ: 8.3.2 156. In orbital hybridization, two or more different atomic orbitals mix to form the same total number of equivalent hybrid orbitals. For instance, the s and p orbitals of an atom combine to make hybrid orbitals having the character of both the s orbital and the p orbital. These hybrid orbitals are equivalent. Orbital hybridization occurs in the methane molecule in which one 2s orbital and three 2p orbitals hybridize to form four sp orbitals. DIF: L3 REF: p. 234 OBJ: 8.3.3 157. ANS: A polar molecule is one in which one end of the molecule has a slightly negative electric charge and the other end has a slightly positive electric charge. An example of a polar molecule is water. The oxygen atom in water develops a slightly negative charge and the hydrogen atoms develop slightly positive charges because of the difference in electronegativity between the oxygen and hydrogen atoms. DIF: L2 REF: p. 239 OBJ: 8.4.1 158. The relative electronegativity of the two bonded atoms determines the polarity of a bond. If the difference in electronegativities between the two atoms is less than 0.4, the bond is nonpolar covalent. If the difference in electronegativities between the two atoms is 0.4 to 1.0, the bond is moderately polar covalent. If the difference in electronegativities between the two atoms is 1.0 to 2.0, the bond is highly polar covalent. If the difference in electronegativities between the two atoms is more than 2.0, the bond is ionic. DIF: L3 REF: p. 237, p. 238, p. 239 OBJ: 8.1.1, 8.4.1

159. Dispersion forces are the weakest of all molecular interactions, and are thought to be caused by the motion of electrons. Generally, the strength of dispersion forces increases as the number of electrons in a molecule increases. Diatomic molecules of halogen elements are an example of molecules whose attraction for one another is caused by dispersion forces. DIF: L2 REF: p. 240 OBJ: 8.4.3 160. Network solids are substances in which all of the atoms are covalently bonded to each other. Melting these substances requires breaking covalent bonds throughout the solid. Two examples are diamond and silicon carbide. DIF: L2 REF: p. 243 OBJ: 8.4.4