Unit C -- Review - Wikispaces--+Unit+2... · Unit C -- Review Multiple Choice ... 48. What are the four quantum numbers in the quantum mechanical model of the atom? What do these

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Unit C -- Review

Multiple ChoiceIdentify the choice that best completes the statement or answers the question.

____ 1. Which object(s) would you use to describe the shape of the 2p orbital?a. a dumb-bellb. a circlec. a sphered. two perpendicular dumb-bellse. a doughnut

____ 2. Which situation must be true for two electrons to occupy the same orbital?a. The electrons must have the same principal quantum number, but the other quantum

numbers must be different.b. The electrons must have the same spin.c. The electrons must have identical sets of quantum numbers.d. The electrons must have low energy.e. The electrons must have the opposite spin.

____ 3. An electron has the following set of quantum numbers: n = 3, l = 1, ml = 1, ms = 1

2. In which orbital is this

electron found?a. 3sb. 3pc. 3dd. 3fe. 4p

____ 4. Which element contains a full 3s orbital?a. Bb. Nac. Mgd. Bee. Ne

____ 5. Which set of quantum numbers is not possible?

a. n = 3, l = 0, ml = 0, ms =1

2

b. n = 5, l = 3, ml = 2, ms =1

2

c. n = 4, l = 3, ml = 1, ms = 1

2

d. n = 5, l = 3, ml = 3, ms = 1

2

e. n = 4, l = 4, ml = 2, ms = 1

2

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____ 6. Which set of quantum numbers is not possible?

a. n = 5, l = 3, ml = 0, ms = 1

2

b. n = 1, l = 0, ml = 0, ms =1

2

c. n = 3, l = 2, ml = 1, ms =1

2

d. n = 4, l = 3, ml = 3, ms =1

2

e. n = 5, l = 2, ml = 0, ms = 1

2

____ 7. Which scientist postulated that electrons can only move between certain energy levels?a. Rutherfordb. Daltonc. Einsteind. Schodingere. Bohr

____ 8. Which element has the ground state electron configuration [Ne] 3s23px13py

1 for its valence electrons?a. Mgb. Alc. Sid. Pe. S

____ 9. Which electron configuration represents a reactive non-metallic element?a. 1s2 2s2 2p6 3s2 3p5

b. 1s2 2s2 2p6 3s2 3p1

c. 1s2 2s2 2p6 3s2

d. 1s2 2s2 2p6 3s2 3p6

e. 1s2 2s2 2p6 3s2 3p6 4s2

____ 10. How many p orbitals are in each energy level, except n = 1?a. 1b. 3c. 5d. 6e. 7

____ 11. What is the maximum number of electrons in n = 3?a. 2b. 3c. 6d. 9e. 18

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____ 12. What is the total number of electrons in the 2p orbitals of a sulfur atom at ground state?a. 8b. 6c. 4d. 3e. 2

____ 13. Which sublevel, when full, corresponds to the first row of transition elements?a. 3db. 3fc. 4dd. 4fe. 4p

____ 14. Which sublevel, when full, corresponds to the lanthanide series of elements?a. 3db. 3fc. 4dd. 4fe. 5f

____ 15. Which pair of atoms and/or ions is isoelectronic?a. O2and Cl

b. Ca2+ and Cl

c. Fand N2

d. Liand Na

e. Kand Kr

____ 16. How does atomic radius change from left to right across a period in the periodic table?a. It increases.b. It decreases.c. It stays the same.d. It increases and then decreases.e. It decreases and then increases.

____ 17. Which atom or ion is isoelectric with Ar?a. Ca2+

b. Kc. S

d. P2

e. Cl

____ 18. Which element has the highest electron affinity?a. Lib. Nc. Od. Fe. Ni

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____ 19. Which metal is the most reactive?a. Alb. Kc. Cud. Zne. Ca

____ 20. Which element has the largest atomic radius?a. Mgb. Bec. Fd. Cle. Si

____ 21. Which element has the smallest effective nuclear charge?a. Alb. Sc. Id. Bee. Na

____ 22. Which element has the lowest first ionization energy?a. Cab. Csc. Brd. Oe. Ba

____ 23. What is the bond angle in a bent molecule, such as water?a. 90b. 104.5c. 107.3d. 109.5e. 120

____ 24. What is the bond angle in a trigonal pyramidal molecule?a. 90b. 104.5c. 107.3d. 109.5e. 120

____ 25. What is the bond angle in a trigonal planar molecule?a. 90b. 104.5c. 107.3d. 109.5e. 120

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____ 26. What is the shape of a molecule of carbon tetrachloride, CCl4?a. linearb. trigonal planarc. trigonal pyramidald. tetrahedrale. angular

____ 27. What is the shape of a molecule of antimony(III) fluoride, SbF3?a. linearb. trigonal planarc. trigonal pyramidald. tetrahedrale. angular

____ 28. What is the shape of a molecule of selenium tetrachloride, SeCl4?a. T-shapedb. seesawc. trigonal bipyramidald. square pyramidale. square planar

____ 29. What is the shape of a molecule of krypton tetraiodide, KrI4?a. T-shapedb. seesawc. trigonal bipyramidald. square pyramidale. square planar

____ 30. Predict the shape of a phosphate ion, PO43.

a. seesawb. trigonal planarc. trigonal pyramidald. tetrahedrale. square planar

____ 31. Predict the shape of OF3+.a. seesawb. trigonal planarc. trigonal pyramidald. tetrahedrale. T-shaped

____ 32. Which compound has polar covalent bonds?a. AgCl2b. CH4

c. Cl2d. CF4

e. B2H8

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____ 33. Which compound is truly covalent?a. SO2

b. MgOc. NH3

d. PCl3e. P2S3

____ 34. Which molecule has a molecular dipole?a. CCl4

b. NF3

c. BeF2

d. CF4

e. CO2

____ 35. Metals can be rolled into sheets and stamped into various forms. In contrast, diamond is very hard and brittle.Which explanation for these different properties is correct?a. Metals have semi-ionic bonds, whereas diamond has covalent bonds.b. The electrons that surround a metal atom are free to move through the metal. The bonding

electron pairs in a diamond are held tightly between two carbon atoms in an overalltetrahedral pattern.

c. The electrons of a metal are held more tightly to the parent atom than the electrons ofcarbon. Hence, the bonds in a metal are stronger than the bonds in diamond.

d. Diamond has strong double bonds between carbon atoms. Metal bonds are normally singlecovalent bonds, which bend easily.

e. Metals are made of metal atoms, whereas diamond is made of non-metal carbon atoms.

____ 36. Which bond is most polar?a. COb. CNc. BOd. BNe. SO

____ 37. Which compound is the best example of a molecular solid held together by London dispersion forces?a. KIb. Na2Oc. H2Od. CO2

e. NH3

____ 38. Which compound is a network solid?a. NaCl (table salt)b. C25H52 (paraffin wax)c. S8 (sulfur)d. mixture of C and Fe (steel)e. SiO2 (quartz)

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____ 39. A pure substance melts at 113C and does not conduct electricity in either the solid or liquid state. How wouldyou classify this substance?a. ionic solidb. network solidc. molecular solidd. metallic solide. atomic solid

____ 40. A solid compound consists of ions bound in a crystal lattice. Which property would you not expect this solid tohave?a. high melting pointb. good conductivity in the solid statec. hardd. soluble in a polar solvente. brittle

____ 41. A substance is a white powder, which melts at 660C to give a transparent liquid. The substance does notconduct electricity in its solid state, but it does conduct electricity in its liquid state. How would you classify thesubstance?a. ionicb. networkc. moleculard. metallice. atomic

____ 42. A substance is a white solid, which melts at a temperature greater than 700C. The substance is extremely hardand not workable. It is a poor conductor of heat and electricity in both its solid and liquid states. How wouldyou classify the substance?a. ionicb. networkc. moleculard. metallice. atomic

____ 43. A substance is a silvery white solid, which melts at 675C to give a silvery liquid. The substance conductselectricity in both its solid and liquid states. How would you classify the substance?a. ionicb. networkc. moleculard. metallice. atomic

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____ 44. A substance has an extremely low boiling point of 4 K. It is a clear, colourless gas at room temperature, and itdoes not conduct electricity in its liquid or solid state. How would you classify this substance?a. ionicb. networkc. moleculard. metallice. atomic

Short Answer

For the following questions, write the most appropriate answer in the space provided.

45. Describe and explain the general trends in atomic radius in the periodic table.

46. What contributions did Schrödinger make to the model of the atom?

47. What contributions did Planck and Einstein make to the current model of the atom?

48. What are the four quantum numbers in the quantum mechanical model of the atom? What do these numbersrepresent?

49. Define the term “orbital.” Describe the shape of each type of orbital. State how many of each type are present,and in which energy levels they are located.

50. Explain what the aufbau principle is and how it is used.

51. Explain Hund’s rule and the Pauli exclusion principle. Give an example to show how these two rules are used.

52. What was one problem that the Bohr model of the atom could not explain?

53. Explain what happens when an electron moves from n = 1 to n = 4. In your explanation, use the followingterms: absorption, emission, photons.

54. What happens when an electron moves from n = 6 to n = 4? In your explanation, use the following terms:absorption, emission, photons.

55. What are the allowed values of ml for an electron with each orbital-shape quantum number.a) l = 3b) l = 1

56. What are the allowed values of l for an electron with each principal quantum number.a) n = 4b) n = 6

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57. What requires more energy: removing a valence electron from an atom or removing an electron from an innerenergy level of an atom? Explain why.

58. Distinguish between the terms “ionization energy” and “electron affinity.”

59. What is the relationship between atomic radius and ionization energy? Explain your answer.

60. What are the allowed values of l and ml if n = 2? What is the total number of orbitals in this energy level?

61. What are the possible values of ml if n = 4 and l = 2? What kind of orbital is described by these quantumnumbers? How many orbitals can be described by these quantum numbers?

62. What contributions did de Broglie make to the understanding of the atom?

63. Distinguish between electron-group arrangement and molecular shape. How does knowing the electron-grouparrangement help you determine the geometric arrangement of the atoms in a molecule?

64. Electron pairs position themselves to minimize the forces of repulsion between them. Put the following pairs inorder, from least to greatest repulsion: lone pair-lone pair, bond pair-bond pair, and lone pair-bond pair.

65. Outline the steps you would take to determine the molecular shape of a compound.

66. Explain why the presence of a bond dipole in a molecule does not necessarily mean there is a molecular dipole.

67. Using water as an example, differentiate between intermolecular forces and intramolecular forces.

68. List the different types of intermolecular forces, and give an example for each.

69. What does VSEPR stand for? How is this theory used to predict molecular shape?

70. What is the difference between dipole-dipole forces and ion-dipole forces?

71. What are allotropes? Give an example.

72. Explain the contributions of Dr. R.J. LeRoy and Dr. R. Bader (two Canadian scientists) to atomic andmolecular theory.

73. What are resonance structures? Give an example of a molecule that requires resonance structures to representits bonding, and draw these resonance structures.

74. CF4 contains polar CF bonds. Is this molecule polar or non-polar?

75. What is a chemical bond? Why do atoms bond

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76. Draw the Lewis structure for SO2. How many bonded pairs of electrons and how many lone pairs of electronsare around the sulfur atom

77. Draw orbital diagrams and Lewis structures to show how each pair of elements can combine. Write thechemical formula of the product.a) Al and Clb) Mg and O

78. What is lattice energyExplain how lattice energy, ionization energy, and electron affinity affect the amount ofenergy that is released when an ionic substance bonds.

79. Explain the difference between covalent bonding and polar covalent bonding, using an example of each.

80. What are the three different molecular shapes that can occur with a tetrahedral geometric arrangement ofelectron groups?

Problem

GraphicsFor the following questions, use the graphics provided to review terms or skills. Add any missing labels,draw any missing parts, or use the graphics to help you answer a question.

81. Using a diagram, explain Rutherford’s gold foil experiment. What did this experiment prove?

82. Complete the following table to describe the different quantum numbers.

Quantum number name Symbol Allowed values Property

83. In the following table, write a complete set of quantum numbers for the last four electrons in an atom that hasthe electron configuration [Ar] 4s2 3d2.

ElectronPrinciple quantum

numberOrbital-shape

quantum numberMagnetic

quantum numberSpin quantum

number1234

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84. Fill in the following outline of a periodic table to show the four energy sublevel (s, p, d, and f) blocks.

85. In the following table, write a complete set of quantum numbers for the last four electrons in an atom that hasthe electron configuration [Ne] 3s2 3p2.


numberOrbital-shape

quantum numberMagnetic quantum

numberSpin quantum

number1234

86. Sketch the shapes of the five different d orbitals on a three-dimensional axis.

87. Sketch the shapes of the three different p orbitals on a three-dimensional axis.

88. On the following outline of a periodic table, use arrows to indicate the directions of increasinga) atomic radiusb) ionization energyc) electron affinity.

89. Consider the following graph of ionization energy.

a) Describe the trend you observe down a group. Explain this trend, using the atomic model.b) Describe the trend you observe across a group. Explain this trend, using the atomic model.

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90. What is hydrogen bonding? Draw a diagram to show the bonding structure of a molecule with hydrogenbonding. (Make sure that you show the bond dipole.) Draw a second diagram to show hydrogen bondingbetween several molecules.

91. Complete the following table to describe the three intramolecular forces.

Force Model/diagram Nature of attraction Energy (kJ/mol) Exampleionic

covalent

metallic

92. Complete the following table to describe the six intermolecular forces.

Force Model/diagram Nature of attraction Energy (kJ/mol) Exampleion-dipole

hydrogen bond

dipole-dipole

dipole-induceddipoleion-induceddipoledispersion

93. Explain the trends in boiling point that are shown in the following graph.

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94. Create a table to summarize the differences in the physical properties of the different types of solids (atomic,molecular, covalent network, ionic, and metallic).

95. What is the free-electron model, as it relates to metallic bonding? How can the free-electron model explain theproperties (such as conductivity, malleability and ductility, and melting and boiling points) of metallic solids?Explain, with the help of diagrams.

96. Draw the three-dimensional representation of each molecule.a) CCl4 (tetrahedral)b) BrCl4(square planar)c) NF3 (trigonal pyramidal)

97. Determine the molecular shape of each molecule, and draw the three-dimensional representation.a) AsCl5

b) H2COc) OCl2

Critical ThinkingFor the following questions, write the answer in the space provided. Use complete sentences in your answer.If the question requires mathematical calculations, show all of your work. Write a final statement that givesyour solution.

98. The ionization energies of a given atom are IE1 = 800 kJ/mol, IE2 = 2400 kJ/mol, IE3 = 3700 kJ/mol, IE4 = 25000 kJ/mol, and IE5 = 32 800 kJ/mol. Predict the valence electron configuration for the atom, and explain yourreasoning.

99. The ionization energies of a given atom are IE1 = 500 kJ/mol, IE2 = 7300 kJ/mol, and IE3 = 11 800 kJ/mol.Predict the valence electron configuration for the atom, and explain your reasoning.

100. Explain what is wrong with each set of quantum numbers.a) n = 3, l = 3, ml = 2; name: 3db) n = 5, l = 3, ml = 5; name: 5f

101. Explain what is wrong with each set of quantum numbers.a) n = 2, l = 1, ml = 0; name: 2pb) n = 3, l = 2, ml = 2; name: 3d

102. Fill in the missing value(s) in each set of quantum numbers.a) n = ?, l = 1, ml = 1; name: 3pb) n = 1, l = ?, ml = ?; name: 1s

103. Fill in the missing values in each set of quantum numbers.a) n = 3, l = 2, ml = ?; name: ?b) n = ?, l = ?, ml = 3; name: 4f

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104. What is the name of the orbital that is represented by each set of quantum numbers?a) n = 2, l = 1, m1 = 1b) n = 4, l = 2, ml = 0

105. What is the name of the orbital that is represented by each set of quantum numbers?a) n = 1, l = 0, m1 = 0b) n = 5, l = 3, ml = 3

106. Write the electron configuration for magnesium.

107. Write the electron configuration for sulfur.

108. Write the electron configuration for each element or ion.a) Mnb) Cac) Ag+

109. Write the electron configuration for each element or ion.a) Pb) Te2

c) U

110. Write the electron configuration for each element or ion.a) Aub) Nic) Be2+

111. Write the condensed electron configuration for each element or ion.a) Bib) Zn2+

c) Al

112. Write the condensed electron configuration for each element or ion.a) Cb) Scc) Zr4+

113. Draw the orbital diagram for each element or ion.a) Cob) Sic) N3

114. Draw the orbital diagram for each element or ion.a) Feb) Na+

c) Cl

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115. Draw the orbital diagram for each element or ion.a) F

b) Arc) Se

116. Identify the group number, period number, and orbital block of the periodic table for the elements whose atomshave the following electron configurations.a) [Ar] 4s2 3d10 4p5

b) [Ne] 3s2

117. Identify the group number, period number, and orbital block of the periodic table for the elements whose atomshave the following electron configurations.a) [Ar] 4s1

b) [Ne] 3s2 3p3

118. Identify four possible elements that have each valence electron configuration.a) s2 p5

b) s2 d3

119. Use the aufbau principle to write the complete electron configuration for an atom of the element that fits eachdescription.a) Group 15 (VA) element in Period 3b) Group 6 (VIB) element in Period 4

120. A hypothetical molecule has the formula XY4. The XY bond is polar covalent, but the molecule is not polar.Which molecular shapes are possible and impossible for the molecule? Explain why, for each shape.

121. The two substances in the following table have similar molecular masses but very different boiling points.Explain how this is possible.

Name Formula Molecular mass Boiling pointethanol CH3CH2OH 46.0 u 78.5°Ccarbon dioxide CO2 44.0 u -78.5°C

122. A 50 mL sample of pure liquid cyclohexane is added to 50 mL of pure liquid methanol. No chemical reactionoccurs. Discuss the intermolecular forces that could exist in this mixture. Predict the end result of mixing thesetwo chemicals.

123. Which substance would you expect to have a higher boiling point: cyclohexane or hexane? Justify your choice.

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124. Consider the properties of the four substances in the following table. Which substance is most likely a metallicsolid? Justify your answer by explaining how the structure of a metal accounts for the properties given.

Substance Boiling point Electrical conductivity in solid and liquid states Other propertiesA 2630°C poor conductor in both hardB 1770°C conductive in liquid state but not in solid state hardC 2220°C conductive in both liquid and solid states thermal conductorD 35°C poor conductor in both soft

125. Consider the properties of the four substances in the following table. Which substance is most likely an ionicsolid? Justify your answer by explaining how the structure of an ionic solid accounts for the properties given.

Substance Boiling point Electrical conductivity in solid and liquidstates

Other properties

A 2630°C poor conductor in both hardB 1770°C conductive in liquid state but not in solid state hardC 2220°C conductive in both liquid and solid states thermal conductorD 35°C poor conductor in both soft

126. Consider the properties of the four substances in the following table. Which substance is most likely a networkcovalent solid? Justify your answer by explaining how the structure of a network solid accounts for theproperties given.


Other properties


127. Consider the properties of the four substances in the following table. Which substance is most likely amolecular solid? Justify your answer by explaining how the structure of a molecular solid accounts for theproperties given.


Other properties


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128. What types of intermolecular forces affect each compound or mixture? Explain your reasoning.a) CH2Cl2b) C5H12

c) O2 dissolved in H2Od) NaCl dissolved in H2O

129. What types of intermolecular forces affect each compound or mixture? Explain your reasoning.a) HFb) C6H12

c) N2 dissolved in H2Od) MgI2

130. Element X has an electronegativity of 2.4, and element Y has an electronegativity of 3.5. The two elementsreact to form a trigonal pyramidal shaped molecule, XY3. Describe and discuss the intermolecular forces thatexist in a sample of this liquid substance.

131. Substances A and B are white crystalline solids. One is an ionic solid, and the other is a molecular solid.Outline three methods that you could use to distinguish each solid.

132. Under what conditions is a tetrahedral shaped molecule also a polar molecule?

133. For each molecule,- draw the Lewis structure- use the VSEPR theory to predict the shape of the compound- decide whether or not the molecule is polar- determine the bond angles- draw the dipole movement for the molecule if the molecule is polara) SCl2b) BF3

c) CHF3

134. For each molecule,- draw the Lewis structure- use the VSEPR theory to predict the shape of the compound- decide whether or not the molecule is polar- determine the bond angles- draw the dipole movement for the molecule if the molecule is polara) CS2

b) PCl3c) SiH4

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135. For each molecule,- draw the Lewis structure- use the VSEPR theory to predict the shape of the compound- decide whether or not the molecule is polar- determine the bond angles- draw the dipole movement for the molecule if the molecule is polara) NH4

+

b) AsCl5c) ICl

136. For each molecule,- draw the Lewis structure- use the VSEPR theory to predict the shape of the compound- decide whether or not the molecule is polar- determine the bond angles- draw the dipole movement for the molecule if the molecule is polara) CH2Cl2b) PH3

c) SeI6

ApplicationsFor the following questions, write the answer in the space provided. Use complete sentences in your answer.If the question requires mathematical calculations, show all of your work. Write a final statement that givesyour solution.

137. Based on what you have learned about atomic emission spectra, explain how the northern lights are created.

138. In 1991, hikers found a man who had been trapped in a glacier. Dating techniques were able to determine thatthe man was 5300 years old. Several tools, including an axe, were found with the man. Scientists determinedthe material that the axe blade was made from, without damaging the axe. Based on information you learned inthis chapter, how did they do this?

139. What intermolecular force is very important in humans? What important function does this force allow tohappen?

140. Suppose that you have a sample of an unknown solid. Outline the steps you would take to identify the type ofsolid it is. Indicate the possible results for each type of solid.

141. Fractional distillation is used to separate the different straight-chain alkanes of crude oil. In fractionaldistillation, a sample of crude oil is heated. The different alkanes become vapours at different temperatures, sothey can be separated. Explain why the different alkanes can be separated in this way.

142. What is a superconductor? Why are superconductors not widely used in technology?

143. KEVLAR is a polymer that is used as a strong, lightweight fabric in products such as canoes and bulletproofvests. How does KEVLAR work?

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144. Describe some uses of amorphous solids.

145. a) Describe the composition of soda-lime glass.b) Glass is an isotropic material. What does this mean, in terms of the properties of glass?

146. Pyrex is a type of glass that is made of 13% B2O3 by weight. What benefit does Pyrex have over regularglass?

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Unit C -- ReviewAnswer Section

MULTIPLE CHOICE

1. ANS: A PTS: 1 DIF: easy REF: K/U2. ANS: E PTS: 1 DIF: easy REF: K/U3. ANS: B PTS: 1 DIF: easy REF: K/U4. ANS: C PTS: 1 DIF: easy REF: K/U5. ANS: E PTS: 1 DIF: average REF: K/U6. ANS: D PTS: 1 DIF: average REF: K/U7. ANS: E PTS: 1 DIF: easy REF: K/U8. ANS: C PTS: 1 DIF: easy REF: K/U9. ANS: A PTS: 1 DIF: average REF: K/U

10. ANS: B PTS: 1 DIF: easy REF: K/U11. ANS: E PTS: 1 DIF: easy REF: K/U12. ANS: C PTS: 1 DIF: average REF: K/U13. ANS: A PTS: 1 DIF: easy REF: K/U14. ANS: D PTS: 1 DIF: easy REF: K/U15. ANS: B PTS: 1 DIF: average REF: K/U16. ANS: B PTS: 1 DIF: easy REF: K/U17. ANS: A PTS: 1 DIF: average REF: K/U18. ANS: D PTS: 1 DIF: average REF: K/U19. ANS: B PTS: 1 DIF: average REF: K/U20. ANS: A PTS: 1 DIF: average REF: K/U21. ANS: E PTS: 1 DIF: average REF: K/U22. ANS: B PTS: 1 DIF: average REF: K/U23. ANS: B PTS: 1 DIF: easy REF: K/U24. ANS: C PTS: 1 DIF: easy REF: K/U25. ANS: E PTS: 1 DIF: easy REF: K/U26. ANS: D PTS: 1 DIF: average REF: K/U | I27. ANS: C PTS: 1 DIF: average REF: K/U | I28. ANS: B PTS: 1 DIF: difficult REF: K/U | I29. ANS: E PTS: 1 DIF: difficult REF: K/U | I30. ANS: D PTS: 1 DIF: difficult REF: K/U | I31. ANS: C PTS: 1 DIF: difficult REF: K/U | I32. ANS: D PTS: 1 DIF: average REF: I33. ANS: E PTS: 1 DIF: average REF: I34. ANS: B PTS: 1 DIF: average REF: I35. ANS: B PTS: 1 DIF: average REF: K/U36. ANS: C PTS: 1 DIF: easy REF: I37. ANS: D PTS: 1 DIF: average REF: K/U | I38. ANS: E PTS: 1 DIF: average REF: K/U | I

ID: A

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39. ANS: C PTS: 1 DIF: difficult REF: I | K/U40. ANS: B PTS: 1 DIF: average REF: K/U | I41. ANS: A PTS: 1 DIF: difficult REF: K/U | I42. ANS: B PTS: 1 DIF: difficult REF: K/U | I43. ANS: D PTS: 1 DIF: difficult REF: K/U | I44. ANS: E PTS: 1 DIF: difficult REF: K/U | I

SHORT ANSWER

45. ANS:Atomic radius increases as you go down a group of elements in the periodic table. This trend is a result ofincreasing numbers of electrons occupying increasing numbers of energy levels. The effective nuclear chargechanges only slightly and therefore does not offset the increase in size due to the increase in energy levels.

Atomic radius decreases as you go left to right across a period in the periodic table. The valence electronsare found in orbitals of the same energy level. At the same time, the effective nuclear charge is increasing withthe increase in nuclear charge, which results in a greater force of attraction pulling the valence electrons closerto the nucleus. Thus, atomic size decreases.

PTS: 1 DIF: average REF: K/U | C46. ANS:

Schrödinger used mathematics and statistics to combine de Broglie’s idea of matter waves and Einstein’s ideaof quantized energy particles. Schrödinger’s mathematical equations and their interpretations, together withHeisenberg’s uncertainty principle, resulted in the birth of the field of quantum mechanics.

PTS: 1 DIF: easy REF: K/U | C47. ANS:

Planck proposed that matter at the atomic level can absorb or emit only discrete quantities of energy. In otherwords, Planck said that the energy of the atom is quantized. Einstein proposed that all forms of electromagneticenergy travel as photons of energy.


1) The principal quantum number, n, indicates the energy level of an atomic orbital and its relative size.2) The orbital-shape quantum number, l, indicates the shape of the orbital.3) The magnetic quantum number, ml, indicates the orientation of the orbital.4) The spin quantum number, ms, indicates the direction in which the electron is spinning.

PTS: 1 DIF: average REF: K/U | C

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49. ANS:1) The s orbital is spherical in shape. One s orbital is present in each energy level.2) The three p orbitals are shaped like dumb-bells. They are present in all energy levels, starting with the

second energy level.3) There are five d orbitals. Four of them have four lobes arranged at right angles to each other, in one plane.

One d orbital has two lobes and a ring of probability in the plane, at right angles to the lobes. The d orbitalsare present in all energy levels higher than the third energy level.

4) The seven f orbitals have eight lobes arranged in three dimensions. They are present in all energy levels afterthe fourth energy level.


The aufbau principle is the imaginary process of building up the ground state electron structure for each atom,in order of atomic number. When determining the electron configuration of an element, the electrons are writtensequentially in orbitals of increasing energy, starting with the electron in the 1s orbital.


The Pauli exclusion principle states that no two electrons in an atom have the same four quantum numbers. (Inother words, no two electrons can occupy the same orbital with the same spin.) For example, boron’s electronconfiguration is 1s2 2s2 2p1.

Hund’s rule states that whenever electrons are added to orbitals of the same energy sub-level, each orbitalreceives one electron before any pairing occurs. When electrons are added singly to separate orbitals of thesame energy sublevel, the electrons must all have the same spin. For example, the electron configuration ofnitrogen is 1s2 2s2 2px

1 2py1 2pz

1.


The Bohr model successfully explained only one-electron systems. It was unable to explain the emission spectrafor atoms with two or more electrons.


When an electron moves from n = 1 to n = 4, the electron absorbs a photon of energy that is equal to thedifference in energy between the two levels.


When an electron moves from n = 6 to n = 4, the electron emits a photon of energy that is equal to thedifference in energy between the two levels.


a) ml = 3, 2, 1, 0 1, 2, 3b) ml = 1, 0, 1

PTS: 1 DIF: easy REF: K/U | I

ID: A

4

56. ANS:a) l = 0, 1, 2, 3b) l = 0, 1, 2, 3, 4, 5

PTS: 1 DIF: easy REF: K/U | I57. ANS:

More energy is required to remove an inner electron because the force of attraction holding this electron isgreater. The greater force of attraction results from the electron being closer to the nucleus. The shielding effectof the inner electrons is reduced, compared with the shielding effect in the valence level, thereby increasing theeffective nuclear charge.


Ionization energy is the energy that is required to completely remove one electron from a ground state gaseousatom. Electron affinity is the change in energy that occurs when an electron is added to a gaseous atom.


Ionization energy is the energy that is required to remove an electron completely from a ground state gaseousatom. This energy tends to increase as the atomic radius decreases. The closer the electrons are to the nucleus,the great the force of attraction pulling or holding the electrons in the atom.


If n = 2, then l = 0, 1. For l = 0, ml = 0. For l = 1, ml = 1, 0, 1. There are four orbitals in this energy level.


The values of ml are 2, 1, 0, 1, 2. The five orbitals described by l = 2 are called the d orbitals.

PTS: 1 DIF: average REF: I62. ANS:

De Broglie conceived the idea that matter has wave-like properties. He developed an equation that enabled himto calculate the wave-length associated with any object, regardless of its size. This equation supported the ideathat electrons have wave-like properties.


Electron-group arrangement refers to the positions of groups of valence electrons around the central atom.Molecular shape refers to the relative positions of the atoms in a molecule. Using the fact that electron pairsrepel each other, the geometric arrangement of the atoms can be predicted by arranging them so that theinteractions of the valence shell electrons of the central atom are minimized. This is the basis of the VSEPRtheory.

PTS: 1 DIF: average REF: K/U

ID: A

5

64. ANS:bond pair-bond pair < lone pair-bond pair < lone pair-lone pair

PTS: 1 DIF: easy REF: K/U65. ANS:

1) Draw a Lewis structure of the molecule.2) Determine the number of pairs of valence electrons around the central atom3) Determine which one of the five geometric arrangements can accommodate this total number of electron

groups.4) Determine the molecular shape from the positions occupied by the bonding pairs and lone pairs.


A molecular dipole results when there are negative and positive regions in a molecule. If the molecule issymmetrical, the individual bond dipoles cancel each other and result in a zero molecular dipole.


Intermolecular forces are forces between molecules. Water has hydrogen bonds, dipole-dipole forces, anddispersion forces that hold adjacent water molecules together. Intramolecular forces are forces within amolecule. They hold the atoms together within the molecule. Each water molecule has polar covalent bondsbetween the hydrogen atoms and the oxygen atom.


Intermolecular forces Exampleion-dipole sodium ions in waterhydrogen bonds waterdipole-dipole iodine monochlorideion-induced dipole ferrous ions and oxygen moleculesdipole-induced dipole hydrochloric acid and chlorinedispersion (London) forces fluorine gas


VSEPR stands for Valence-Shell Electron-Pair Repulsion. The VSEPR theory is used to predict the shapes ofmolecules by relating the total number of electron groups around the central atom to a geometric arrangementthat minimizes the repulsions of these electron groups. For example, two electron groups arrange themselves ina linear position, with bond angles of 180.


Dipole-dipole forces occur between molecules that have molecular dipoles: for example, between ICl molecules.Ion-dipoles occur between the ions of an ionic solid and a molecule with a molecular dipole: for example,between Na+ and H2O.


ID: A

6

71. ANS:Allotropes are different crystalline or molecular forms of the same element that have different physical andchemical properties. Carbon is an example of an element that has different allotropes. Two of its allotropes aregraphite and diamond.


Dr. R.J. LeRoy, in conjunction with Dr. R. Bernstein, developed the Le Roy-Bernstein theory. This theorydescribes the patterns of energies, and other properties, for the vibrational levels of molecules that are veryclose to dissociation. Le Roy then developed a calculation for determining the Le Roy radius: the minimumdistance at which the equation describing the vibrations of molecules close to dissociation is thought to be valid.

Dr. R. Bader invented the theory of “Atoms in Molecules” (the AIM theory). This theory links themathematics of quantum mechanics to the atoms and bonds in a molecule. It adopts electron density, which isrelated to the Schrödinger description of the atom, as a starting point for mapping molecules.


Resonance structures are models that give the same relative position of atoms as Lewis structures, but showdifferent places for their bonding and lone pairs. An example is SO2.


The bond between C and F is a polar covalent bond because CF4 is tetrahedral in shape. The effects of the bonddipoles cancel each other, however, resulting in a non-polar molecule.


Chemical bonds are electrostatic forces that hold atoms together in compounds. In nature, systems of lowerenergy tend to be favoured over systems of higher energy. Bonded atoms tend to have lower energy than single,uncombined atoms.


There are three bonded pairs of electrons and one lone pair of electrons around the sulfur atom.

PTS: 1 DIF: easy REF: I

ID: A

7

77. ANS:a)

The product is AlCl3.b)

The product is MgO.

PTS: 1 DIF: easy REF: I78. ANS:

Lattice energy is the amount of energy that is released when an ionic crystal forms from the gaseous ions of itselements. The ionization energy of the metal is always greater than the electron affinity of the non-metal. Thus,the transfer of electrons between gaseous ions requires energy. The lattice energy that is released favours theformation of ionic bonds.


ID: A

8

79. ANS:Covalent bonding exists when the bonding electrons are shared equally or nearly equally, as in molecules of O2

and N2. Polar covalent bonding exists when there is unequal sharing of a pair of electrons between two atoms,as in HCl and H2O.


tetrahedral, trigonal pyramidal, and angular (bent)

PTS: 1 DIF: easy REF: K/U

PROBLEM

81. ANS:

This experiment proved that the atom is made up mainly of empty space, with a small, massive region ofconcentrated charge at the centre. The charge was soon determined to be positive.


Quantum number name Symbol Allowed values Propertyprincipal n 1, 2, 3, 4… energy levelorbital-shape l 0, 1, 2 . . . (n - 1) shape of the orbitalmagnetic ml -l, . . . -1, 0, 1, . . . +l orientation of the orbitalspin ms

1

2,

1

2direction of electron spin

PTS: 1 DIF: easy REF: K/U

ID: A

9

83. ANS:


numberOrbital-shape


numberSpin quantum

number1 4 0 0

1

2

2 4 0 0 1

2

3 3 2 -2 1

2

4 3 2 -1 1

2




numberOrbital-shape


numberSpin quantum

number1 3 0 0

1

2

2 3 0 0 1

2

3 3 1 -1 1

2

4 3 1 0 1

2

PTS: 1 DIF: average REF: K/U86. ANS:


ID: A

10

87. ANS:



a) Ionization energy decreases as you go down a group. This is caused by electrons occupying higher energylevels and therefore being farther from the nucleus. The increase in distance makes it easier to remove anelectron completely from an atom. Thus, ionization energy decreases.

b) As you go left to right across a period, ionization energy increases. This is caused by increasing effectivenuclear charge, resulting from increasing nuclear charge from left to right.

PTS: 1 DIF: average REF: I | C90. ANS:

Hydrogen bonding is a particularly strong form of dipole-dipole attraction. It exists between a hydrogen atom ina polar-bonded molecule that contains HN, HO, or HF bonds and an unshared pair of electrons in asmall, electronegative N, O, or F atom of an adjacent molecule.


ID: A

11

91. ANS:

PTS: 1 DIF: average REF: K/U92. ANS:


ID: A

12

93. ANS:With the exception of H2O, HF, and NH3, the trends in boiling point are consistent with an increase in thenumber of electrons and therefore an increase in dispersion forces. Hence, the hydrids of Group 14 (IVA)elements have the lowest boiling points, while the hydrids of Group 17 (VIIA) elements have the highest boilingpoints. The boiling points increase with increasing number of electrons in the hydrids of each group. The threeexceptionsH2O, HF, and NH3have much higher boiling points than the trends observed because of stronghydrogen bonding forces.


Type of CrystalSolid

Boiling Point Electrical Conductivityin Liquid State

Other Physical Properties ofCrystals

Atom low very low very softMolecular generally low

(non-polar);intermediatepolar

very low non-polar: very soft; soluble innon-polar solventspolar: somewhat hard, but brittle;many are soluble in water

CovalentNetwork

very high low hard crystals that are insoluble inmost liquids

Ionic high high hard and brittle; many dissolve inwater

Metallic most high very high all have a lustre, are malleable andductile, and are good conductors;they dissolve in other metals toform alloys


ID: A

13

95. ANS:The free electron model shows a metal composed of a densely packed core of metallic cations, within adelocalized region of shared, mobile valence electrons.

Conductivity: Metals are good conductors of heat and electricity because electrons can move freely throughoutthe metallic structure.Malleability and ductility: The malleability of metals can be explained by thinking of metallic bonds asnon-directional. The positive ions are often layered as a fixed array. When stress is applied to a metal, one layerof positive ions can slide over another layer. The layers move without breaking the array.

Melting and boiling points: The melting and boiling points of the Group 1 (IA) metals are generally lower thanthe melting and boiling points of the Group 2 (IIA) metals. This is because the Group 1 (IA) elements have onlyone s electron and smaller nuclear charge, while the Group 2 (IIA) elements have two s electrons, larger nuclearcharge, and smaller size. The transition elements have very high melting and boiling points, since the delectrons and the electrons of the valence shell have similar magnitudes. Therefore, the d electrons canparticipate in the metallic bonding.


ID: A

14

96. ANS:



The atom probably has three valence electrons because the increase from the third to the fourth ionizationenergies is much greater than the increase from the first to the second to the third ionization energies. This largeincrease can be explained by the fact that the fourth electron is being removed from a filled energy level, whichis closer to the nucleus. Thus, the outer (valence) electrons have all been removed.

PTS: 1 DIF: difficult REF: I | C99. ANS:

The atom probably has one valence electron because the second ionization energy is much greater than the firstionization energy. This large increase can be explained by the fact that the second and third electrons are beingremoved from a filled energy level, which is closer to the nucleus. Thus, the outer (valence) electrons have allbeen removed.


a) l cannot equal 3, since the maximum value of l is (n 1) and d orbitals have an l value of 2.b) ml cannot equal 5, since the maximum value of ml is +l.


a) l cannot equal 1, since values of l can only be positive integers.b) This set of quantum numbers is a valid set.

PTS: 1 DIF: average REF: I

ID: A

15

102. ANS:a) n = 3b) l = 0, ml = 0


a) ml = 2, 1, 0, 1, or 2; name: 3db) n = 4, l = 3


a) 2pb) 4d


a) 1sb) 5f


1s2 2s2 2p6 3s2


1s2 2s2 2p6 3s2 3p4


a) 1s2 2s2 2p6 3s2 3p6 4s2 3d5

b) 1s2 2s2 2p6 3s2 3p6 4s2

c) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d8


a) 1s2 2s2 2p6 3s2 3p3

b) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6

c) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f4


a) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d9

b) 1s2 2s2 2p6 3s2 3p6 4s2 3d8

c) 1s2


ID: A

16

111. ANS:a) [Xe] 6s2 4f14 5d1 6p3

b) [Ar] 4s2 3d10

c) [Ne] 3s2 3p1


a) [He] 2s2 2p2

b) [Ar] 4s2 3d1

c) [Kr] or [Ar]4s2 3d10 4p6




ID: A

17

115. ANS:


a) Group VIIA (17), Period 4, p blockb) Group IIA (2), Period 2, s block


a) Group IA (1), Period 4, s blockb) Group VA (15), Period 3, p block


a) The configuration s2p5 corresponds to Group VIIA (17). Fluorine, chlorine, bromine, and iodine fit thisconfiguration.

b) The configuration s2d3 corresponds to Group VB (5). Vanadium, niobium, tantalum, and dubnium fit thisconfiguration.


a) 1s2 2s2 2p6 3s2 3p5

b) 1s2 2s2 2p6 3s2 3p6 4s2 3d4

PTS: 1 DIF: difficult REF: I120. ANS:

Possible shapes are tetrahedral and square planar. These two shapes are symmetrical. Thus, the individual bonddipoles cancel out, leaving the molecule without a molecular dipole.

An impossible shape is seesaw. This shape is not symmetrical, so the individual bond dipoles do not cancelout.

PTS: 1 DIF: difficult REF: I

ID: A

18

121. ANS:The CO2 molecule has a linear shape, with the oxygen atoms located on either side of the carbon atom. Thisresults in a non-polar molecule. The only forces that are holding the CO2 molecules together are dispersionforces. CH3CH2OH has oxygen atoms bonded to hydrogen atoms, resulting in hydrogen bonds betweenmolecules. These hydrogen bonds are much stronger than the dispersion forces between CO2 molecules.Therefore, CH3CH2OH has a much higher boiling point.


Cyclohexane is a symmetrical molecule with no dipole. The forces that are holding together the cyclohexanemolecules are dispersion forces. Methanol has polarity and hydrogen bonding. The two compounds are notmiscible, and two layers will appear in the container. There will be no decrease in total volume.


The chain molecule, hexane, has a higher boiling point. The atoms are more spread out (have more surfacearea) in this molecule, and hence more chances to form intermolecular interactions.


Substance C is a metallic solid because it is an electrical conductor in both the solid and liquid states. It is alsoa thermal conductor. The conductivity of substance C can be explained by the delocalized valence electrons,which are mobile and free to move throughout the array of positive ions.


Substance B is most likely an ionic solid. Ionic solids conduct in the liquid state but do not conduct in the solidstate. As a solid, the ions are fixed in position and unable to move. When the solid is liquefied, the ions are freeto move, allowing the movement of charge through the liquid.


Substance A is most likely a network solid. Due to the presence of covalent bonds between the atoms, networksolids do not conduct in either the liquid or solid state. They are extremely hard and have very high meltingpoints.


Substance D is most likely a molecular solid. Molecular solids are held together by weaker intermolecularforces, such as dispersion forces, dipole-dipole forces, and hydrogen bonding. These weaker forces result inlower boiling points and softness. The bonds in molecular solids are covalent or polar covalent, so the electronsare not free to move throughout the solid. As a result, molecular solids cannot conduct electricity in either thesolid or liquid state.

PTS: 1 DIF: difficult REF: I | C

ID: A

19

128. ANS:a) Dipole-dipole: The molecules of CH2Cl2 have a molecular dipole.b) Dispersion forces: The C5H12 (hexane) molecule has covalent bonds between atoms.c) Dipole-induced dipole: The H2O molecule has a bent shape, so it has a molecular dipole. This dipole induces

a dipole in the O2 molecules, resulting in their mutual attraction.d) Ion-dipole interactions: NaCl dissociates into its ions as it dissolves in H2O. The H2O molecule has a

molecular dipole. Hence, ion-dipole interaction takes place between the ions of NaCl and the polar H2Omolecules.


a) Hydrogen bonds: The bonding of hydrogen to the small electronegative fluorine atom results in the chargedistribution that allows the formation of hydrogen bonds between molecules.

b) Dispersion forces: The C6H12 molecule has covalent bonds between atoms.c) Dipole-induced dipole: The H2O molecule has a bent shape, so it has a molecular dipole. This dipole induces

a dipole in the N2 molecules, resulting in their mutual attraction.d) Ion-dipole interactions. MgI2 dissociates into its ions as it dissolves in H2O. The H2O molecule has a

molecular dipole. Hence, ion-dipole interaction takes place between the ions of MgI2 and the polar H2Omolecules.


Since the electronegativity difference between element X and element Y is 1.1, there is polar covalent bondingand a bond dipole on each bond. Since the molecular shape is trigonal pyramidal, the molecule is notsymmetrical and therefore has a molecular dipole. The intermolecular forces that exist between molecules ofthis substance in the liquid state are therefore dipole-dipole interactions, due to the molecular dipole anddispersion forces.


1) Place a small sample of each substance in a deflagrating spoon. Heat both samples in a Bunsen burnerflame. The molecular solid will melt very quickly. The ionic solid will not melt, or it will melt slowly usinga hot flame on the Bunsen burner.

2) Dissolve a small amount of each solid in a sample of water. Using conductivity apparatus, test theconductivity of each solution. The solution that contains the ionic solid will conduct. The solution thatcontains the molecular solid will not conduct.

3) Apply a force to a sample of each solid. The ionic solid will be brittle. The molecular solid will be soft andtherefore deform more readily.


A tetrahedral molecular is also a polar molecule if the individual bonds have dipoles and the atoms surroundingthe central atom are not the same. The individual bond dipoles do not cancel each other, so the molecule has amolecular dipole.

PTS: 1 DIF: average REF: I | C

ID: A

20

133. ANS:


PTS: 1 DIF: difficult REF: I

ID: A

21

135. ANS:



Scientists believe that the northern lights are caused by streams of protons and electrons (plasma) emanatingfrom the Sun. When the electrons from the plasma interact with gaseous atoms in the Earth’s upperatmosphere, the atoms emit coloured light.

PTS: 1 DIF: average REF: MC | C

ID: A

22

138. ANS:The axe was irradiated with high-energy X ray radiation. The atoms in the axe absorbed the radiation, causingelectrons from a lower energy level to be ejected from the atom. This caused electrons from an outer energylevel to “move in,” to occupy the vacated space. As the electrons fell to a less energetic state, they emitted Xrays. The electrons of each atom emitted X rays of a particular wavelength. Scientists used this energyspectrum “signature” to identify the atoms.

PTS: 1 DIF: average REF: MC | C139. ANS:

Hydrogen bonding is very important in humans. Hydrogen bonds allow the formation of the spiral-like, doublehelix structure of DNA (deoxyribonucleic acid). As well, they are responsible for the properties of water.

PTS: 1 DIF: easy REF: MC140. ANS:

1) Heat a small sample of the unknown solid in a crucible, in a Bunsen burner flame. If the solid readily melts,then it is likely a molecular solid. If the solid melts only after prolonged heating, then it is possibly ionic ormetallic. If it does not melt, then it is possibly a network solid.

2) Using conductivity apparatus, determine if a sample of the solid conducts electricity. If it does conductelectricity, then the unknown is most likely metallic. If it does not conduct electricity, the unknown might bea molecular, ionic, or network solid.

3) Melt a small sample of the solid. Test for electrical conductivity again. If the liquid conducts electricity, andif the solid also conducted electricity, then the solid is metallic. If the liquid conducts electricity, but the soliddid not, then the solid is ionic. If the liquid does not conduct electricity, then the solid is either molecular ornetwork.

4) If the unknown solid did not conduct electricity as either a solid or a liquid, and if it had a low melting point,then it is a molecular substance. If the unknown did not conduct electricity as a solid or a liquid, and if it hada high melting point, then it is a network solid.


As the length of the alkane chain increases, so do the mass and the surface area between the molecules. Theincrease in surface area results in an increase in dispersion forces. The greater the dispersion forces, the greaterthe boiling point is. Thus, the small alkane molecules vaporize first, at lower temperatures, and the largeralkanes vaporize at higher temperatures. Thus, the alkanes can be separated according to their molecular size.

PTS: 1 DIF: average REF: K/U | C | MC142. ANS:

A superconductor is a material that has no resistance to electrical current. Superconductors are not widely usedin technology due to the temperatures at which they function. At the present time, the highest temperature atwhich a known material will act as a superconductor is 138 K.

PTS: 1 DIF: average REF: MC

ID: A

23

143. ANS:KEVLARbelongs to a class of compounds called aramids, which are polymers that contain aromatic andamide groups. The strength of the polymer comes from the intramolecular forces of the aromatic group. Theseintramolecular forces limit bond rotation in the straight chains, as well as the amide linkage that leads tointermolecular hydrogen bonding between the straight chains. The strands are twisted to increase their densityand thickness. The strands are then woven very tightly. A resin coating is applied and sandwiched between twolayers of plastic film. Each layer acts as a net and slows down a puncture or a bullet, a little at a time, until ithas been stopped.

PTS: 1 DIF: average REF: C | MC144. ANS:

Solids that are not arranged in orderly, crystalline structures are called amorphous solids. Melting silica andadding limestone and soda ash produces glass. Soda-lime glass is used to make window and bottle glass. 90%of all the glass that is produced is soda-lime glass. Adding boron oxide allows Pyrex to be made. Flint glassis used to make lenses, due to its low refractive index. Liquidmetal alloy is an amorphous alloy of zirconium,which is used to make golf clubs.

PTS: 1 DIF: average REF: MC145. ANS:

a) Soda-lime glass is an amorphous solid that is formed using molten silica, which contains Na2O, CaO, MgO,and Al2O3.

b) An isotropic material is a material that has the same properties in all directions. Therefore, glass expandsuniformily as it is heated.

PTS: 1 DIF: average REF: MC146. ANS:

Pyrex has a very low thermal expansion coefficient, which means that it expands very little when heated.Pyrex can be heated numerous times over a flame without cracking.

PTS: 1 DIF: average REF: MC

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Unit C -- Review - Wikispaces--+Unit+2... · Unit C -- Review Multiple Choice ... 48. What are the four quantum numbers in the quantum mechanical model of the atom? What do these