Vsepr and lewis dot Matching



Download 235,7 Kb.
Date conversion11.09.2018
Size235,7 Kb.
vsepr and lewis dot
Matching
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

____ 1. a depiction of the arrangement of atoms in molecules and polyatomic ions


____ 2. a covalent bond in which only one pair of electrons is shared
____ 3. a covalent bond in which two pairs of electrons are shared
____ 4. a covalent bond in which the shared electron pair comes from only one of the atoms
____ 5. a covalent bond between two atoms of significantly different electronegativities
____ 6. 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

____ 7. energy needed to break a single bond between two covalently bonded atoms


____ 8. symmetrical bond along the axis between the two nuclei
____ 9. molecular orbital that can be occupied by two electrons of a covalent bond
____ 10. 109.5
____ 11. shapes adjust so valence-electron pairs are as far apart as possible
____ 12. attraction between polar molecules
____ 13. crystal in which all the atoms are covalently bonded to each other
Multiple Choice

Identify the choice that best completes the statement or answers the question.
____ 14. 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.

____ 15. 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

____ 16. Which of these elements does not exist as a diatomic molecule?



a.

Ne

c.

H

b.

F

d.

I

____ 17. 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.

____ 18. 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

____ 19. Which of the following elements can form diatomic molecules held together by triple covalent bonds?



a.

carbon

c.

fluorine

b.

oxygen

d.

nitrogen

____ 20. Which noble gas has the same electron configuration as the oxygen in a water molecule?



a.

helium

c.

argon

b.

neon

d.

xenon

____ 21. 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

____ 22. 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

____ 23. 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.

2s2p

c.

2s2p3s

b.

2s2p3s

d.

2s2p

____ 24. Which of the following diatomic molecules is joined by a double covalent bond?



a.



c.



b.



d.



____ 25. A molecule with a single covalent bond is ____.



a.

CO

c.

CO

b.

Cl

d.

N

____ 26. 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

____ 27. 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.

____ 28. When H forms a bond with HO to form the hydronium ion HO, 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

____ 29. Which of the following bonds is the least reactive?



a.

C—C

c.

O—H

b.

H—H

d.

H—Cl

____ 30. 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

____ 31. In which of the following compounds is the octet expanded to include 12 electrons?



a.

HS

c.

PCl

b.

PCl

d.

SF

____ 32. How many electrons can occupy a single molecular orbital?



a.

0

c.

2

b.

1

d.

4

____ 33. 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

____ 34. 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

____ 35. 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

____ 36. 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

____ 37. 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

____ 38. 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

____ 39. 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

____ 40. 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

____ 41. The shape of the methane molecule is called ____.



a.

tetrahedral

c.

four-cornered

b.

square

d.

planar

____ 42. 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

____ 43. 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

____ 44. Experimental evidence suggests that the H—C—H bond angles in ethene, CH, are ____.



a.

90

c.

120

b.

109.5

d.

180

____ 45. What type of hybrid orbital exists in the methane molecule?



a.

sp

c.

sp

b.

sp

d.

spd

____ 46. What is the shape of a molecule with a triple bond?



a.

tetrahedral

c.

bent

b.

pyramidal

d.

linear

____ 47. 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.



____ 48. How many pi bonds are formed when sp hybridization occurs in ethene, CH?



a.

0

c.

2

b.

1

d.

3

____ 49. Which of the following atoms acquires the most negative charge in a covalent bond with hydrogen?



a.

C

c.

O

b.

Na

d.

S

____ 50. 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

____ 51. Which of the following covalent bonds is the most polar?



a.

H—F

c.

H—H

b.

H—C

d.

H—N

____ 52. 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

____ 53. 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

____ 54. Which of the forces of molecular attraction is the weakest?



a.

dipole interaction

c.

hydrogen bond

b.

dispersion

d.

single covalent bond

____ 55. 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

____ 56. What are the weakest attractions between molecules?



a.

ionic forces

c.

covalent forces

b.

Van der Waals forces

d.

hydrogen forces

____ 57. 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

____ 58. 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.

____ 59. Which type of solid has the highest melting point?



a.

ionic solid

c.

metal

b.

network solid

d.

nonmetallic solid

____ 60. 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



Numeric Response
61. How many valence electrons does an iodine atom have?
62. What is the total number of covalent bonds normally associated with a single carbon atom in a compound?
63. How many electrons are shared in a single covalent bond?
64. How many electrons does a nitrogen atom need to gain in order to attain a noble-gas electron configuration?
65. How many unshared pairs of electrons does the nitrogen atom in ammonia possess?
66. How many electrons does carbon need to gain in order to obtain a noble-gas electron configuration?
67. How many electrons are shared in a double covalent bond?
68. How many covalent bonds are in a covalently bonded molecule containing 1 phosphorus atom and 3 chlorine atoms?
69. How many unshared pairs of electrons are in a molecule of hydrogen iodide?
70. What is the bond angle in a water molecule?
Essay
71. What is bond dissociation energy, and how does it affect carbon compounds?
72. Can some atoms exceed the limits of the octet rule in bonding? If so, give an example.
73. Indicate how bonding is explained in terms of molecular orbitals.
74. Explain a pi bond and a sigma bond. Which of these bond types tends to be the weaker? Why?
75. Explain what is meant by VSEPR theory. Give an example of how VSEPR theory can be applied to predict the shape of a molecule.
76. Explain what is meant by orbital hybridization. Give an example of a molecule in which orbital hybridization occurs.
77. Explain what a polar molecule is. Provide an example.
78. What determines the degree of polarity in a bond? Distinguish between nonpolar covalent, polar covalent, and ionic bonds in terms of relative polarity.
79. 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.
80. Describe a network solid and give two examples.

vsepr and lewis dot

Answer Section
MATCHING
1. ANS: C PTS: 1 DIF: L1 REF: p. 218

OBJ: 8.1.2 Describe the information a molecular formula provides.


2. ANS: D PTS: 1 DIF: L1 REF: p. 217

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e


3. ANS: B PTS: 1 DIF: L1 REF: p. 221

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e


4. ANS: A PTS: 1 DIF: L1 REF: p. 223

OBJ: 8.2.4 Distinguish between a covalent bond and a coordinate covalent bond and describe how the strength of a covalent bond is related to its bond dissociation energy.

STA: SC3.e
5. ANS: E PTS: 1 DIF: L1 REF: p. 238

OBJ: 8.4.1 Describe how electronegativity values determine the charge distribution in a polar molecule.


6. ANS: F PTS: 1 DIF: L1 REF: p. 241

OBJ: 8.4.3 Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds.


7. ANS: D PTS: 1 DIF: L1 REF: p. 226

OBJ: 8.2.5 Describe how oxygen atoms are bonded in ozone. STA: SC3.e


8. ANS: G PTS: 1 DIF: L1 REF: p. 230

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


9. ANS: B PTS: 1 DIF: L1 REF: p. 230

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


10. ANS: E PTS: 1 DIF: L1 REF: p. 232

OBJ: 8.3.2 Describe how VSEPR theory helps predict the shapes of molecules.


11. ANS: F PTS: 1 DIF: L1 REF: p. 232

OBJ: 8.3.2 Describe how VSEPR theory helps predict the shapes of molecules.


12. ANS: C PTS: 1 DIF: L1 REF: p. 240

OBJ: 8.4.3 Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds.


13. ANS: A PTS: 1 DIF: L1 REF: p. 243

OBJ: 8.4.4 Identify the reason network solids have high melting points.


MULTIPLE CHOICE
14. ANS: D PTS: 1 DIF: L2 REF: p. 244

OBJ: 8.1.1 Distinguish between the melting points and boiling points of molecular compounds and ionic compounds. STA: SC3.e


15. ANS: A PTS: 1 DIF: L1 REF: p. 215

OBJ: 8.1.2 Describe the information a molecular formula provides.


16. ANS: A PTS: 1 DIF: L1 REF: p. 217

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e


17. ANS: C PTS: 1 DIF: L2 REF: p. 217

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e


18. ANS: B PTS: 1 DIF: L2 REF: p. 217

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e


19. ANS: D PTS: 1 DIF: L2 REF: p. 221

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e


20. ANS: B PTS: 1 DIF: L2 REF: p. 218

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e


21. ANS: D PTS: 1 DIF: L3 REF: p. 217 | p. 218

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e


22. ANS: D PTS: 1 DIF: L1 REF: p. 218

OBJ: 8.2.2 Demonstrate how electron dot structures represent shared electrons.

STA: SC3.e
23. ANS: D PTS: 1 DIF: L3 REF: p. 220 | p. 234

OBJ: 8.2.2 Demonstrate how electron dot structures represent shared electrons.

STA: SC3.e
24. ANS: A PTS: 1 DIF: L2 REF: p. 221

OBJ: 8.2.3 Describe how atoms form double or triple covalent bonds.

STA: SC3.e
25. ANS: B PTS: 1 DIF: L2 REF: p. 222

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. | 8.2.4 Distinguish between a covalent bond and a coordinate covalent bond and describe how the strength of a covalent bond is related to its bond dissociation energy.

STA: SC3.e
26. ANS: C PTS: 1 DIF: L2 REF: p. 223

OBJ: 8.2.4 Distinguish between a covalent bond and a coordinate covalent bond and describe how the strength of a covalent bond is related to its bond dissociation energy.

STA: SC3.e
27. ANS: D PTS: 1 DIF: L2 REF: p. 223

OBJ: 8.2.4 Distinguish between a covalent bond and a coordinate covalent bond and describe how the strength of a covalent bond is related to its bond dissociation energy.

STA: SC3.e
28. ANS: A PTS: 1 DIF: L2 REF: p. 225

OBJ: 8.2.4 Distinguish between a covalent bond and a coordinate covalent bond and describe how the strength of a covalent bond is related to its bond dissociation energy.

STA: SC3.e
29. ANS: B PTS: 1 DIF: L2 REF: p. 226

OBJ: 8.2.5 Describe how oxygen atoms are bonded in ozone. STA: SC3.e


30. ANS: A PTS: 1 DIF: L1 REF: p. 227

OBJ: 8.2.5 Describe how oxygen atoms are bonded in ozone. STA: SC3.e


31. ANS: D PTS: 1 DIF: L2 REF: p. 229

OBJ: 8.2.5 Describe how oxygen atoms are bonded in ozone. STA: SC3.e


32. ANS: C PTS: 1 DIF: L1 REF: p. 230

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


33. ANS: C PTS: 1 DIF: L1 REF: p. 230

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


34. ANS: C PTS: 1 DIF: L1 REF: p. 230

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


35. ANS: D PTS: 1 DIF: L2 REF: p. 230

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


36. ANS: C PTS: 1 DIF: L2 REF: p. 231

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


37. ANS: B PTS: 1 DIF: L2 REF: p. 231

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


38. ANS: D PTS: 1 DIF: L2 REF: p. 230 | p. 231

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


39. ANS: D PTS: 1 DIF: L2 REF: p. 231

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


40. ANS: A PTS: 1 DIF: L1 REF: p. 232

OBJ: 8.3.2 Describe how VSEPR theory helps predict the shapes of molecules.


41. ANS: A PTS: 1 DIF: L1 REF: p. 232

OBJ: 8.3.2 Describe how VSEPR theory helps predict the shapes of molecules.


42. ANS: A PTS: 1 DIF: L2 REF: p. 233

OBJ: 8.3.2 Describe how VSEPR theory helps predict the shapes of molecules.


43. ANS: C PTS: 1 DIF: L1 REF: p. 234

OBJ: 8.3.3 Identify the ways in which orbital hybridization is useful in describing molecules.


44. ANS: C PTS: 1 DIF: L1 REF: p. 235

OBJ: 8.3.3 Identify the ways in which orbital hybridization is useful in describing molecules.


45. ANS: C PTS: 1 DIF: L2 REF: p. 234

OBJ: 8.3.3 Identify the ways in which orbital hybridization is useful in describing molecules.


46. ANS: D PTS: 1 DIF: L2 REF: p. 235

OBJ: 8.3.3 Identify the ways in which orbital hybridization is useful in describing molecules.


47. ANS: C PTS: 1 DIF: L2 REF: p. 235

OBJ: 8.3.3 Identify the ways in which orbital hybridization is useful in describing molecules.


48. ANS: B PTS: 1 DIF: L2 REF: p. 235

OBJ: 8.3.3 Identify the ways in which orbital hybridization is useful in describing molecules.


49. ANS: C PTS: 1 DIF: L2 REF: p. 238 | p. 239

OBJ: 8.4.1 Describe how electronegativity values determine the charge distribution in a polar molecule.


50. ANS: C PTS: 1 DIF: L2 REF: p. 238 | p. 239

OBJ: 8.1.1 Distinguish between the melting points and boiling points of molecular compounds and ionic compounds. | 8.4.1 Describe how electronegativity values determine the charge distribution in a polar molecule. STA: SC3.e


51. ANS: A PTS: 1 DIF: L3 REF: p. 238 | p. 239

OBJ: 8.4.1 Describe how electronegativity values determine the charge distribution in a polar molecule.


52. ANS: A PTS: 1 DIF: L3 REF: p. 239

OBJ: 8.4.2 Describe what happens to polar molecules when they are placed between oppositely charged metal plates.


53. ANS: B PTS: 1 DIF: L1 REF: p. 240

OBJ: 8.4.3 Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds.


54. ANS: B PTS: 1 DIF: L1 REF: p. 240

OBJ: 8.4.3 Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds.


55. ANS: B PTS: 1 DIF: L1 REF: p. 240

OBJ: 8.1.1 Distinguish between the melting points and boiling points of molecular compounds and ionic compounds. | 8.4.3 Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds. STA: SC3.e


56. ANS: B PTS: 1 DIF: L1 REF: p. 240

OBJ: 8.4.3 Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds.


57. ANS: D PTS: 1 DIF: L2 REF: p. 241

OBJ: 8.4.3 Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds.


58. ANS: A PTS: 1 DIF: L2 REF: p. 241

OBJ: 8.4.1 Describe how electronegativity values determine the charge distribution in a polar molecule. | 8.4.3 Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds.


59. ANS: B PTS: 1 DIF: L1 REF: p. 243

OBJ: 8.4.4 Identify the reason network solids have high melting points.


60. ANS: D PTS: 1 DIF: L1 REF: p. 243

OBJ: 8.4.4 Identify the reason network solids have high melting points.


NUMERIC RESPONSE
61. ANS: 7

PTS: 1 DIF: L1 REF: p. 218

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e
62. ANS: 4

PTS: 1 DIF: L1 REF: p. 219

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e
63. ANS: 2

PTS: 1 DIF: L2 REF: p. 217

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e
64. ANS: 3

PTS: 1 DIF: L2 REF: p. 219

OBJ: 8.2.1 Describe how electrons are shared to form a covalent bonds and identify exceptions to the octet rule. STA: SC3.e
65. ANS: 1

PTS: 1 DIF: L2 REF: p. 219

OBJ: 8.2.2 Demonstrate how electron dot structures represent shared electrons.

STA: SC3.e


66. ANS: 4

PTS: 1 DIF: L2 REF: p. 219

OBJ: 8.2.2 Demonstrate how electron dot structures represent shared electrons.

STA: SC3.e


67. ANS: 4

PTS: 1 DIF: L2 REF: p. 221

OBJ: 8.2.2 Demonstrate how electron dot structures represent shared electrons.

STA: SC3.e


68. ANS: 3

PTS: 1 DIF: L3 REF: p. 219

OBJ: 8.2.2 Demonstrate how electron dot structures represent shared electrons.

STA: SC3.e


69. ANS: 3

PTS: 1 DIF: L3 REF: p. 220

OBJ: 8.2.2 Demonstrate how electron dot structures represent shared electrons.

STA: SC3.e


70. ANS: 105

PTS: 1 DIF: L2 REF: p. 233

OBJ: 8.3.2 Describe how VSEPR theory helps predict the shapes of molecules.
ESSAY
71. ANS:

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.

PTS: 1 DIF: L2 REF: p. 226

OBJ: 8.2.5 Describe how oxygen atoms are bonded in ozone. STA: SC3.e


72. ANS:

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.

PTS: 1 DIF: L2 REF: p. 228 | p. 229

OBJ: 8.2.5 Describe how oxygen atoms are bonded in ozone. STA: SC3.e


73. ANS:

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.

PTS: 1 DIF: L3 REF: p. 230

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


74. ANS:

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.

PTS: 1 DIF: L3 REF: p. 230 | p. 231

OBJ: 8.3.1 Describe the relationship between atomic and molecular orbitals.


75. ANS:

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.

PTS: 1 DIF: L3 REF: p. 232

OBJ: 8.3.2 Describe how VSEPR theory helps predict the shapes of molecules.


76. ANS:

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.

PTS: 1 DIF: L3 REF: p. 234

OBJ: 8.3.3 Identify the ways in which orbital hybridization is useful in describing molecules.


77. 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.

PTS: 1 DIF: L2 REF: p. 239

OBJ: 8.4.1 Describe how electronegativity values determine the charge distribution in a polar molecule.


78. ANS:

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.

PTS: 1 DIF: L3 REF: p. 237 | p. 238 | p. 239

OBJ: 8.1.1 Distinguish between the melting points and boiling points of molecular compounds and ionic compounds. | 8.4.1 Describe how electronegativity values determine the charge distribution in a polar molecule. STA: SC3.e


79. ANS:

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.

PTS: 1 DIF: L2 REF: p. 240

OBJ: 8.4.3 Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds.


80. ANS:

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.



PTS: 1 DIF: L2 REF: p. 243

OBJ: 8.4.4 Identify the reason network solids have high melting points.


The database is protected by copyright ©sckool.org 2016
send message

    Main page