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Chapter Eight
Bonding:
General Concepts
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Contents
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8-1 Types of Chemical Bonds
Coulomb’
s law
The energy of interaction between a pair of ions
can be calculated using Coulomb’
s law:
E (2.3110 19 J nm)(
Q1Q2
)
r
where E has units joules, r is the distance
between the ion centers in nanometers, and
Q1 and Q2 are the numerical ion charges.
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For example, the distance between the centers of
the Na+ and Cl- ions is 0.276 nm, and the ionic
energy pair of ions is
E (2.3110 19 J nm) [
(1)(1)
] 8.37 10 19 J
0.276 nm
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Questions to Consider

What is meant by the term “
chemical
bond?”

Why do atoms bond with each other to
form molecules?
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Figure 8.1 (a)
The interaction of two
hydrogen atoms.
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Bond
length
Figure 8.1(b) Energy profiles as a function between the
hydrogen atoms. As the atoms approach each other (right side
of graph), the energy decreases until the distances reaches
0.074 nm and then begins to increase again due to repulsions.
Key ideas in bonding

Ionic Bonding: Electrons are transferred

Covalent Bonding: Electrons are shared
equally

What about intermediate cases?

Polar covalent bond:
H
F
 
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R
1
t
c
a
e
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Chemical bond

What is meant by the term “
chemical bond?”

Why do atoms bond with each other to form
molecules?

How do atoms bond with each other to form
molecules?
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8-2 Electronegativity
Linus Pauling(1901-1995)
Expected H-X bonding energy
= ½ (H-H bond energy + X-X bond energy)
△ = (H-X)act - (H-X)exp
If X has a greater electronegativity than H, the shares electron(s)
will tend to be closer to the X atom. The molecule will be polar,
with charge distribution.
The greater is the difference in the electronegativities of the atoms,
the greater is the ionic compound and the greater is the value of △.
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Electronegativity: the ability of an atom in a
molecule to attract shared electrons to itself.
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Figure 8.2 The effect of an electric field on hydrogen fluoride molecules
Polar molecules
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The Pauling electronegativity values
Figure 8-3
The Pauling electronegativity values. Electronegativity generally
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increases across a period and decreases down a group.
2
t
c
a
Re
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The general trend for electronegativity

What is the general trend for
electronegativity across rows and down
columns on the periodic table?

Explain the trend.
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Ex 8.1 Relative Bond Polarities
Order the following bonds according to polarity:
H–H, O–H, Cl–H, S–H, and F–H.
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Table 8.1 The Relationship Between
Electronegativity and Bond Type
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8-3 Bond polarity and dipole moments
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Ex 8.2 Bond Polarity and Dipole
Moment
For each of the following molecules, show the direction
of the bond polarities and indicate which ones have a
dipole moment: HCl, Cl2, SO3(a planar molecule with the
oxygen atoms spaced evenly around the central sulfur
atom), CH4 [trtrahedral(see Table 8.2) with the carbon
atom at the center], and H2S (V-shaped with the sulfur
atom at the point).
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Solution:
(a) HCl
(b) Cl2
(c) SO3
(d) CH4
(e) H2S
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8-4 Ions: Electron
configurations and sizes
Predicting Formulas of Ionic Compounds
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Sizes of ions
Table 8.3 Common ions with noble gas configuration in
ionic compounds
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t
c
a
Re
Choose an alkali metal, an alkaline metal, a noble gas,
and a halogen so that they constitute an isoelectronic
series when the metals and halogen are written as
their most stable ions.

What is the electron configuration for each species?

Determine the number of electrons for each species.

Determine the number of protons for each species.

Rank the species according to increasing radius.

Rank the species according to increasing ionization
energy.
Ionic radii
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What we can “read”from the
periodic table:

Trends for







Atomic size
Ion radius
Ionization energy
Electronegativity
Electron configurations
Predicting formulas for ionic
compounds
Ranking polarity of bonds
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Ex 8.4 Relative Lon Size II
Choose the largest ion in each of the following
groups.
a.Li+, Na+, K+, Rb+, Cs+
b.Ba2+, Cs+, I-, Te2-
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Ex 8.3 Relative Lon Size I
Arrange the ions Se2-, Br-, Rb+, and Sr2+ in order of
decreasing size.
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8-5 Energy effects in binary
ionic compounds
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Formation of an ionic solid
1. Sublimation of the solid metal
• M(s)
 M(g)
[endothermic] (For
Li(s) is +161 kJ.)
2. Ionization of the metal atoms
• M(g)
 M+(g) + e[endothermic] (For Li(g) is +520 kJ)
3. Dissociation of the nonmetal
• 1/2X2(g)
 X(g)
[endothermic] (For F is +½ (154 kJ)
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Formation of an ionic solid (continued)
4. Formation of Xions in the gas phase:
X(g) + e X(g) [exothermic]
(For F- is -328 kJ/mole)
5. Formation of the solid MX:
M+(g) + X(g)  MX(s) [quite exothermic]
(Corresponding to the lattice energy for LiF, which is -1047
kJ./mole)
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Figure 8.11
Comparing
energy changes
Born-Haber cycle for NaCl
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Lattice Energy Calculations
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8.6 Partial ionic character of
covalent bonds
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The relationship between the ionic character
of a covalent bond and the electronegativity
difference of the bonded atoms.
Figure 8.13
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8-7 The Covalent Chemical Bond:
A Model
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Models
Models are attempts to explain how nature
operates on the microscopic level based on
experiences in the macroscopic world.
The Localized Electron Bonding Model
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Fundamental Properties of Models
1.
A model does not equal reality.
2.
Models are oversimplifications, and are
therefore often wrong.
3.
Models become more complicated as they
age.
4.
We must understand the underlying
assumptions in a model so that we don’
t
misuse it.
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8-8 Covalent Bond Energies and
Chemical Reactions
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Ex 8.5 △H from Bond Energies
Using the bond energies listed in Table 8.4, calculate △H
for the reaction of methane with chlorine and fluorine to
give Freon-12(CF2Cl2).
CH 4 ( g ) 2Cl2 ( g ) 2 F2 ( g )  CF2Cl2 ( g ) 2 HF ( g ) 2 HCl ( g )
8-9 The Localized Electron
Bonding Model
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A molecule is composed of atoms that are
bound together by sharing pairs of electrons
using the atomic orbitals of the bound atoms.
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Localized Electron Model
1.
Description of valence electron
arrangement (Lewis structure).
2.
Prediction of geometry (VSEPR model).
3.
Description of atomic orbital types used to
share electrons or hold long pairs.
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8-10 Lewis Structure

Shows how valence electrons are arranged
among atoms in a molecule.

Reflects central idea that stability of a
compound relates to noble gas electron
configuration.
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Lewis Structures
1.
Sum the valence electrons.
2.
Place bonding electrons between pairs of
atoms.
3.
Atoms usually have noble gas
configurations.
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Ex 8.6 Writing Lewis Structures
Give the Lewis structure for each of the following.
a. HF, b. N2, c. NH3, d. CH4, e. CF4, f. NO+
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8-11 Exceptions to the Octet Rule
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Ex 8.7 Lewis Structures for Molecules
That Violate the Octet Rule I
Write the Lewis structure for PCl5.
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Ex 8.8 Lewis Structures for Molecules
That Violate the Octet Rule II
Write the Lewis structure for each molecule or ion.
a. ClF3 b. XeO3 c. RnCl2 d. BeCl2 e. ICl4-
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8-12 Resonance
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Ex 8.9 Resonance Structures
Describe the electron arrangement in the nitrite
anion (NO2-) using the localized electron model.
Rules Governing Formal Charge
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Ex 8.10 Formal Charges
Give possible Lewis structures for XeO3 , an explosive
compound of xenon. Which Lewis structure or structures
are most appropriate according to the formal charges?
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8-13 Molecular Structure:
The VSEPR Model
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VSEPR Model
The structure around a given atom is
determined principally by minimizing
electron pair repulsions.
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Predicting a VSEPR Structure
1.
Draw Lewis structure.
2.
Put pairs as far apart as possible.
3.
Determine positions of atoms from the way
electron pairs are shared.
4.
Determine the name of molecular structure
from positions of the atoms.
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Ex 8.11 Prediction of Molecular
Structure I
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Describe the molecular structure of the water molecule.
Solution
The Lewis structure for water is
There are four pairs of electrons: two bonding pairs and two
nonbonding pairs. To minimize repulsions, these best
arrangement in a tetrahedral array, as shown in Fig. 8.17.
Figure 8.17
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Ex 8.12 Prediction of Molecular Structure II
When phosphorus reacts with excess chlorine gas, the
compound phosphorus pentachloride (PCl5) is formed. In
the gaseous and liquid states, this substance consists of
PCl5 molecules, but in the solid state it consists of a 1 : 1
mixture of PCl4+ and PCl6- ions. Predict the geometric
structures of PCl5, PCl4+, and PCl6-.
Solution
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The Lewis structure for PCl5 is shown. Five pairs of electrons around
the phosphorous atom require a trigonal bipyramidal arrangement
(see Table 8.6).
The Lewis structure for the PCl4+ ions (5+4(7) -1 = 32 valence
electrons) is shown. There are four pairs of electrons surrounding
the phosphorus atom in the PCl4+ ion, which requires a tetrahedral
arrangement of the pairs.
The Lewis structure for PCl6- (5 + 6(7) + 1 = 48 valence electrons)
is shown. Since each electron pair is shared with a chlorine atom, an
octahedral PCl6- anion is predicted.
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Ex 8.13 Prediction of Molecular
Structure III
Because the noble gases have filled s and p valence orbitals,
they were not expected to be chemically reactive. In fact,
for many years these elements were called insert gases
because of this supposed inability to form any compounds.
However. In the early 1960s several compounds of krypton,
xenon, and radon were synthesized. For example, a team at
the Argonne National Laboratory produced the stable
colorless compound xenon tetrafluoride (XeF4). Predict its
structure and whether it has a dipole moment.
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Solution
The Lewis structure for XeF4 is
The xenon atom in this molecule is surrounded by
six pairs of electrons, which means an octahedral
arrangement.
The arrangement in Fig. 8.20(b) is preferred, and the
molecular structure is predicted to be square planar. There
is an octahedral arrangement of electron pairs, but the
atoms form a square planar structure. Although each Xe-F
bond is polar, their structure causes the polarities to cancel.
Thus XeF4 has no dipole moment as shown in the margin.
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Ex 8.14 Structures of Molecules with
Multiple Bonds
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Predict the molecular structure of the sulfur dioxide
molecule. Is this molecule expected to have a dipole moment?
Solution
We must determine the Lewis structure for the SO2
molecule, which has 18 valence electrons. The expected
resonance structures are
The structure of the SO2 molecule expected to be Vshaped, with a 120-degree bond angle. The molecule has a
dipole moment as shown:
Molecules Containing No Single Central Atom
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The VSEPR Model- How Well Does It Work?
VSEPR
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VSEPR: Two Electron Pairs
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VSEPR: Three Electron Pairs
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VSEPR: Four Electron Pairs
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VSEPR: Iodine Pentafluoride
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