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Unit 5 - Molecular Bonding and Structure
Accelerated Chemistry I
Properties of Ionic and Metallic Compounds
1. Write the formula for the ion formed when each of the following elements loses or gains valence
electrons:
a. aluminum
c. nitrogen
b. strontium
d. iodine
2. Which of the following pairs of elements will not form ionic compounds?
a. sulfur and oxygen
b. sodium and calcium
c. sodium and sulfur
d. oxygen and chlorine
3. Using bonding theory, briefly explain why metals are good conductors of heat and electricity.
4. Using bonding theory, briefly explain why ionic compounds are brittle.
5. Name the three crystal arrangements of closely packed metal atoms.
6. Is it accurate to describe sodium chloride (NaCl) as consisting of individual particles, each
composed of one Na+ cation and one Cl- anion? Explain your answer.
LEWIS STRUCTURES FOR IONS AND MOLECULES
Bonding theory explains many of the features observed in chemical bonds. In particular, bonding
theory explains the law of definite composition, provides a basis for predicting molecular structure,
and accounts for observed differences in bond strength, or bond energy.
The idea of the covalent bond was first suggested by the American chemist G. N. Lewis in 1916. He
pointed out that the electron configuration of the noble gases appears to be a particularly stable one.
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Lewis suggested that atoms, by sharing electrons to form an electron-pair bond, can acquire a stable
noble-gas structure.
For the following elements, write out the electron configuration of a ground state atom and draw the
electron dot symbol that represents the valence electrons. Finally, predict how many electrons the
atom needs to gain or lose to become isoelectronic with a noble gas element.
Table 1. Lewis Electron Dot Structures of Atoms and Electrons Needed
Element
Electron Configuration
Electron Dot Symbol
Electrons to Gain
H
C
N
O
F
S
P
Writing Lewis Structures: Lewis Dot Structures of Molecules
For each compound determine the number of valence electrons for each atom in the compound and
them compute the total number of valence electrons. Draw the proper Lewis Dot Structure. Check
your structure using formal charges if necessary.
Formula
Valence Electrons for
Each Atom
Total # of Valence
Electrons
Lewis Dot Structure
1. H2S
2. NH3
3. CH2Cl2
4. H2O2
2
Formula
Valence Electrons for
Each Atom
Total # of Valence
Electrons
Lewis Dot Structure
5. HOCl
6. OH-
7. H3O+
8. NH4+
9. HCN
10. O2
11. N2
12. CO2
13. C2H4
14. C2H2
15. H2CO
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molecular shape of H2OThe water molecule is “bent”
because, focusing on the atoms,
the
107.5°
three atoms are arrange on a bent
Exceptions to the Octet Rule
with#aofbond
angle of 105°.
Valence Electrons for line,
Total
Valence
Formula
Lewis Dot Structure
Each Atom
Electrons
1. NO2
four electron
domains
104.5°
2. BH3
3. BeH2
MOLECULAR SHAPES
The names for molecular shapes are based on the position of the atoms in the molecule—not on the
position of the electron domains. For example, the water molecule H2O has four domains of electrons,
with the orbitals pointing toward the four corners of a tetrahedron. However, we say the
Five common molecular shapes:
Molecular Shapes
1. Explain how the shape of a molecule can be predicted from its Lewis structure.
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2. Why is the bond angle in bent molecules expected to be 104.5°?
3. Draw the Lewis structure for PH3. How many electron domains does the phosphorus atom have in
PH3? What is the electron pair geometry? Name the shape of the PH3 molecule and give the
approximate bond angles.
4. Draw the Lewis structure for SO2. How many electron domains does the sulfur atom have in SO2?
What is the electron pair geometry of this molecule? Name the molecular shape of the SO2
molecule and give the approximate bond angles.
5. Draw the Lewis structure, predict the molecular shape, and give the bond angles for:
a. Cl2
d. C2Cl4
b. CO
e. NO3-
c. SO42-
f. SF3+
6. Explain why the ∠HCC = 180° in HCCH.
7. Explain why the ∠ClNN = 117° (close to 120°) in ClNNCl.
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A
B
Hybridization, Geometry, Sigma and Pi Bonds
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2
C
H H O
1.
hybridization of the central atom in each of the following.
(Hint: Draw the dot
| the
A | Give
B|| expected
C
1
structure
for
the
molecule
first.)
H⎯C⎯C⎯C⎯H
1
a. H2S
2
d. SO2
g. H2CO
A
1
2
b. CO32-
e. CO2
B h. HCNC
c. CCl4
f. AsH3
i. P2H4
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2. Give all the hybridization and bond angles in the following molecules.
a. Acrolein, a component of photochemical smog, has a pungent odor and irritates eyes. What are
the hybridizations of carbons 1 and 2? What are the approximate angles A, B, and C? How many
σ-bonds and how many π-bonds are there?
b. Lactic acid is a natural compound found in sour milk. How many σbonds and how many π-bonds are there? Identify the hybridization of
atoms 1-3. What are the approximate bond angles A thru C?
c.
Cinnamaldehyde occurs naturally in cinnamon oil. How many σbonds and how many π-bonds are there? Identify the hybridization
of carbon atoms 1-3. What are the approximate bond angles A thru
C?
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Resonance
1.
Explain what is meant by the term resonance structure.
2.
Draw all of the resonance structures of:
a. O3
b. CO32-
Molecular Shape and Polarity
1. Define the following terms.
a. dipole
b. dipole moment
2. Explain why an atom cannot have a permanent dipole moment.
3. The C-O bonds in carbon dioxide are polar, yet the dipole moment of the molecule is zero. Explain.
4. What does the dipole moment of a molecule tell us about its intermolecular forces?
5. Draw the dot structure for each of the following molecules. Indicate the shape of each molecule
(linear, trigonal planar, trigonal pyramidal, tetrahedral). Then refer to the electronegativities of the
atoms involved and determine whether each bond is polar. Using lines to represent bonds and dots
to represent nonbonding pairs, draw a representation of the molecule that resembles the molecular
shape. If it is polar, add an arrowhead to the lines representing bonds, pointing toward the more
electronegative atom. Finally, decide whether the molecule as a whole is polar. If it is, draw a large
arrow near the molecule to indicate the direction of polarity. The first substance has been done as
an example.
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••
N
I
I
I
Formula
NI3
Lewis Structure
Shape
Trigonal
pryamidal
Polarity of
Bonds
N-I polar
Representation
Polarity
of
Molecule
Polar
PCl3
BCl3
CO2
CCl4
CH2F2
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6. Indicate whether the following molecules are polar or nonpolar.
a. CO
c. CF4
b. H2O
d. CH2Cl2
7. List these molecules in order of increasing dipole moment: H2O, CBr4, H2S, HF, NH3, CO2.
Write T for true and F for false. If a statement is false, replace the underlined word or phrase with one
that will make the statement true, and write your correction on the blank provided.
8. Dispersion forces are the only attractive forces that attract between nonpolar molecules.
9. Molecules cannot exhibit both dipole and dispersion interactions.
10. Substances composed of nonpolar molecules are generally gases at room temperature or highboiling liquids.
11. Substances composed of polar molecules generally have higher boiling points than do nonpolar
compounds.
12. A nonpolar molecule can have polar bonds.
13. Water is a nonpolar molecule that has polar bonds.
14. The compounds Br2 and ICl have the same number of electrons, yet Br2 melts at -7.2°C whereas
ICl melts at 27.2°C. Explain.
15. Lists these attractive forces in order of strength, starting with the weakest: dispersion forces,
hydrogen bond, covalent bond.
16. List the types of intermolecular forces that exist in each of these species. You may have to draw
the dot structure in order to determine the polarity of the molecule.
(a) benzene (C6H6)
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H O
N C C
H 3Cl H O–H
(b) CH
H
HHHH
H
H
(c) PF3
(d) NaCl
(e) CS2
17. What kind of attractive forces must be overcome to (a) melt ice, (b) boil molecular bromine, (c)
melt solid iodine, and (d) dissociate F2 into F atoms?
18. The table below shows the boiling point of four
compounds. Which compound has the strongest
intermolecular forces? How do you know?
Compound
HF(l)
CH3Cl(l)
CH3F(l)
HCl(l)
Boiling Point (°C)
19.4
-24.2
-78.8
-83.7
19. Of the following, which has the lowest boiling point and why: CH4, CH3CH3, or CH3CH2CH3?
20. Amino acids are the building blocks of proteins. A molecule of the simplest amino acid, glycine, is
shown here. Draw another glycine molecule showing how it may hydrogen bond to another
glycine molecule.
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Electronegativities
2.20
H
0.98
Li
0.93
Na
0.82
K
0.82
Rb
1.57
Be
1.31
Mg
1.00
Ca
0.95
Sr
1.36
Sc
1.22
Y
1.54 1.63 1.66
Ti V Cr
1.33 1.60 2.16
Zr Nb Mo
1.55
Mn
1.9
Tc
1.83
Fe
2.2
Ru
1.88
Co
2.28
Rh
1.91
Ni
2.20
Pd
1.90
Cu
1.93
Ag
1.65
Zn
1.69
Cd
2.04
B
1.61
Al
1.81
Ga
1.78
In
2.55
C
1.90
Si
2.01
Ge
1.96
Sn
3.04
N
2.19
P
2.18
As
2.05
Sb
3.44
O
2.58
S
2.55
Se
2.10
Te
3.98
F
3.16
Cl
2.96
Br
2.66
I
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