Chem 1 Review Sheet
Metallic Bonds & Chapter 9
Dr. Casagrande / Mrs. Rinaldi
Be able to:
1. Describe what a metallic bond is.
2. Describe the electron sea model, delocalized electrons, and how the two are related and how it relates to
properties of metals
3. Describe why a metal deforms when an external force is applied rather than shattering, as ionic
compounds do.
4. Describe a covalent bond and how it differs from an ionic bond.
5. Describe the octet rule and how it applies to formation of covalent bonds.
6. Describe sigma (σ) and pi (π) bonds.
7. Indicate how many pi bonds are there in a double covalent bond or in a triple covalent bond.
8. Connect C & H atoms in a given arrangement with single, double, or triple bonds and indicate how
many σ and π bonds are present.
9. Indicate how does the bond strength relates to the bond length. Rank single, double, and triple bonds in
terms of bond strength and bond length.
10. Describe bond dissociation energy and how it relates to bond strength.
11. Determine the names of covalent compounds and acids from their formulas and their formulas from
their names (remember water—H2O, ammonia—NH3, and methane—CH4)
12. Apply the process of counting valence electrons and placing electron pairs on the structure for drawing
Lewis structures of covalent compounds from their formulas
13. Indicate the bond angles for tetrahedral, trigonal planar, and linear molecules.
14. Draw a 3-D model and indicate the overall geometry, molecular shape, and bond angle for covalent
molecules and polyatomic ions from their formulas
15. Use electronegativity to determine the polar character of a bond.
16. Use the bond polarity and symmetry of a molecule to determine if it is polar.
17. Problems pg. 272-273 #72, 73, 75, 78, 79, 80, 81, 84, 85, 86, 88, 89, 90, 91 (draw the Lewis Structures),
92, 94, 95, 96, 97, 98, 105 (no hybrid orbitals), 108, 109, 110, 111
18. Convert the following formulas to names and names to formulas:
a. N2O
Dinitrogen monoxide
a. carbon dioxide
CO2
b. PF3
Phosphorous trifluoride
b. dinitrogen pentoxide
N2 O5
c. CCl4
Carbon tetrachloride
c. diboron hexahydride
B2H6
d. P2O5
Diphosphorous pentoxide
d. silicon tetrachloride
SiCl4
e. SF6
Sulfur hexafluoride
e. chlorine trifluoride
ClF3
f. ClF5
Chlorine pentafluoride
f. xenon tetrafluoride
XeF4
g. H2Te
hydrotelluric acid
g. acetic acid
HC2H3O2
h. HBrO
hypobromous acid
h. hydroiodic acid
HI
19. For the following bonds, determine ∆EN and polarity for each bond. If the bond is polar covalent,
indicate the dipole using either the arrow or the δ symbols. If ionic, put in the charges.
a.
b.
H—F
∆EN= 4.0 – 2.1 = 1.9
B—N
∆EN= 3.0 – 2.0 = 1.0
Polarity: Polar Covalent
Polarity: Polar Covalent
c.
d.
Ba—I
∆EN= 2.5 – 0.9 = 1.6
C—S
Polarity: Polar Covalent
∆EN= 2.5 – 2.5 = 0
Polarity: Nonpolar Covalent
20. Draw the Lewis structures for the following molecules or polyatomic ions. If the molecule does not have
an expanded octet, determine the electron and molecular geometries and draw the 3-dimensional
structure. On the 3-D drawing, put a dipole arrow (
) or the partial charges (δ+/δ–) on all polar
bonds and indicate whether the overall molecule is polar.
Molecule/
Ion
a. SeO3
Lewis
Structure
Electron
Geometry
24 e–
3-D Drawing
∆ENSe-O=1.1
Se
O
O
O
Trigonal
Planar
O
b. SO2
Molec.
Geometry
Trigonal
Planar
Se
O
18 e–
O
Polar?
No:
Polar
bonds
but not
asymm.
∆ENS-O=1.0
O
S
O
Trigonal
Planar
Bent
O
S
Yes
O
c. SiF4
32 e–
∆ENSi–F=2.2
F
F
Si
F
F
Tetrahedral
Tetrahedral
F
d. PBr3
P
Br
Tetrahedral
Trigonal
Pyramid
8 e–
P
Br
Br
Br
Yes
∆ENH-Te=0
H
Tetrahedral
Te
H
26 e–
C
Cl
Te
Bent
H
H
∆ENH-Cl=0.9; ∆ENC-H=0.4
H
Cl
2
F
∆ENP-Br=0.7
Br
f. CHCl3
F
F
26 e–
Br
e. H2Te
Si
No
(Symmetric)
H
Cl
Tetrahedral
Tetrahedral
C
Cl
Cl
Cl
No:
No Polar
Bonds
Yes:
Asymm
with
polar
bonds
Chapter 9 Review
Problems pg. 272-273 #72, 73, 75, 78, 79, 80, 81, 84, 85, 86, 88, 89, 90, 91 (draw the Lewis Structures), 92,
94, 95, 96, 97, 98, 105 (no hybrid orbitals), 108, 109, 110, 111
72. What is the octet rule, and how is it used in covalent bonding?
Atoms lose, gain, or share electrons to end with a full outer energy level, which contains an octet. In
covalent bonding, nonmetal atoms share electrons to achieve an octet.
73. Describe the formation of a covalent bond.
The nucleus of one atom attracts the electrons of the other atom, and they share one or more pairs of
electrons.
75. Describe the forces, both attractive and repulsive, that occur as two atoms
come closer together.
Electrons are attracted to the nuclei of the atoms, and the nuclei repel each
other.
78. When is a molecular compound named as an acid?
When it starts with H or ends in COOH and releases H+ in water.
79. What must be known in order to draw the Lewis structure for a molecule?
The number of valence electrons for each atom and how the atoms are arranged in the molecule.
80. On what is the VSEPR model based?
The repulsive nature of electron pairs around a central atom.
81. What is the molecular shape of each of the following molecules? Estimate the bond angle for each
assuming no lone pairs.
B linear, 180°;
B
A linear, 180°
a. A
b. A
A
B
A Trigonal planar, 120°;
c. A
d.
Tetrahedral, 109.5°
A
A
B
A
A
84. Describe the trends in electronegativity in the periodic table.
Electronegativity increases moving from bottom to top in a group and left to right in a period.
85. Explain the difference between nonpolar molecules and polar molecules.
A nonpolar molecule has a symmetric distribution of charge, while a nonpolar molecule has an uneven
distribution of electrons with partial positive charge one side of the molecule and partial negative on the
other.
86. Compare the location of bonding electrons in a polar covalent bond with those in a nonpolar covalent
bond. Explain your answer.
Electrons in a polar bond are closer to the more electronegative atom because of unequal sharing. Those
in a nonpolar bond are shared equally.
88. Give the number of valence electrons in N, As, Br, and Se. Predict the number of covalent bonds needed
for each of these elements to satisfy the octet rule.
N: 5 valence electrons, 3 covalent bond; As: 5, 3; Br: 7, 1; Se: 6, 2
Chapter 9 Review
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89. Locate the sigma and pi bonds in the following molecule:
O
H
C
Single bonds: sigma bonds; double bond: one sigma bond, one pi bond.
H
90. Locate the sigma and pi bonds in the following molecule:
H
C
C H
Single bonds: sigma bonds; triple bond: one sigma bond, two pi bonds.
91. Consider the molecules CO, CO2, and CH2O. Which C—O bond is shortest? In which molecule is the
C—O bond strongest?
The bond in CO is a triple bond, while the others are double bonds, so it is shortest and strongest.
92. Consider the carbon-nitrogen bonds in the following:
C
and
N
H
H
H
C
N
H
H
Which bond is shorter? Which is stronger? The triple bond is shorter and stronger.
94. Name each of the following solutions as an acid.
a. HClO2: chlorous acid c. H2Se: hydroselenic acid
b. H3PO4: phosphoric acid
d. HClO3: chloric acid
95. Name each of the following molecules.
a. NF3: nitrogen trifluoride
c. SO3: sulfur trioxide
b. NO: nitrogen monoxide
d. SiF4: silicon tetrafluoride
96. Name each of the following molecules:
a. SeO2: selenium dioxide
c. N2F4: dinitrogen tetrafluoride
b. SeO3: selenium trioxide
d. S4N4: tetrasulfur tetranitride
97. Write the formula for each of the following:
a. sulfur difluoride: SF2
c. carbon tetrafluoride: CF4
b. silicon tetrachloride: SiCl4
d. sulfurous acid: H2SO3
98. Write the formula for each of the following:
a. silicon dioxide: SiO2
c. chlorine trifluoride: ClF3
b. bromous acid: HBrO2
d. hydrobromic acid: HBr
105. Predict the molecular shape and bond angle, and identify the hybrid orbitals for each of the
following. Drawing the Lewis structure may help you.
a. SCl2
c. HOF
b. NH2Cl
d. BF3
a.
Tetrahedral/Bent; <109.5°
S
Cl
Cl
b.
Tetrahedral/Trigonal Pyramidal, <109.5°
N
Cl
H
H
4
Chapter 9 Review
c.
Tetrahedral/Bent; <109.5°
O
F
H
d.
Trigonal planar/Trigonal planar, 120°
F
B
F
F
108. For each pair, indicate the more polar bond by circling the negative end of its dipole.
a. C–S, C–O
b. C–F, C–N
c. P–H, P–Cl
109. For each of the bonds listed, tell which atom is more negatively charged.
a. C–H
C
c. C–S
S
b. C–N
N
d. C–O
O
110. Predict which of the following bonds is the most polar.
a. C–O
c. C–Cl
Si–O has the largest ∆EN, so it is the most polar
b. Si–O
d. C–Br
111. Rank the following bonds according to increasing polarity.
a. C–H ΔEN = 2.5 – 2.1 = 0.4 d. O–H ΔEN = 3.5 – 2.1 = 1.4
b. N–H ΔEN = 3.0 – 2.1 = 0.9 e. Cl–H ΔEN = 3.0 – 2.1 = 0.9
c. Si–H ΔEN = 2.1 – 1.8 = 0.3
Chapter 9 Review
c<a<b=e<d
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