Chemistry
FIFTH EDITION
by Steven S. Zumdahl
University of Illinois
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1
Chemistry
FIFTH EDITION
Chapter 8
Chemical Foundations
Molecular Bonding and Structure play the central role
in determining the course of all chemical reactions.
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Section 8.1
Types of Chemical Bonds
•
Forces that hold groups of
atoms together and make them
function as a unit.
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Bond Energy
 It
is the energy required to break a bond.
 It
gives us information about the strength
of a bonding interaction.
 Various
types of Bonds
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Ionic Bonds
 Formed
from electrostatic attractions of closely
packed, oppositely charged ions.
 Formed
when an atom that easily loses
electrons reacts with one that has a high
electron affinity.
 Usually
when a metal loses electrons &
the electrons are gained by a non-metal.
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Ionic Bonds
To calculate the Energy of Interaction between
a pair of ions -- Coulombs Law.
E = 2.31  10
19
J nm (Q1Q2 / r )
•
Q1 and Q2 = numerical ion charges
•
r = distance between ion centers (in nm)
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Problem
• Calculate the energy of interaction between
Ag+ and Br- if the internuclear distance of
AgBr is 0.120 nm.
• E = 2.31 x 10-19 J nm (Q1 Q2/ r)
• E = 2.31 x 10-19 J nm ((+1) (-1)/ 0.120 nm)
• E = -1.925 x 10-18 J per unit
= -1.93 x 10-18 J per unit
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Problem
• E = -1.925 x 10-18 J per unit
x 6.022 x 1023 units per mole
x 1 kJ / 1000 J
• E = -1159.235 kJ/mole = -1160 kJ/mole
• E = Negative means Attractive force
• Ion Pair has lower energy than the separated
ions.
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8
Covalent Bonding
• Bonding between atoms where electrons are
shared by nuclei.
• A bond will form if the energy of the system is
lower than that of the separated atoms.
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Figure 8.1
Interaction of Two H Atoms and the Energy Profile
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Bond Length
•
The distance where the system
energy is a minimum.
• Corresponds to distance apart where
the combination of repulsive forces
(proton-proton & electron-electron)
and attractive forces (electron-proton)
allows the system to reach a minimum
energy.
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Non-polar Covalent Bond:
Equal Sharing of Electrons
Ex: Br---Br
Polar Covalent Bond:
Unequal Sharing of Electrons
Ex: H---F
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The Effect of an
Electric Field on
Hydrogen Fluoride
Molecules
Molecules tend to line up
With
- ends toward + pole
&
+ ends toward – pole.
Fluorine is slightly neg (δ-)
since it is more
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Section 8.2
Electronegativity
•
The ability of an atom in a molecule
to attract shared electrons to itself.
• Electronegativity Ranges
from 4.0 (Fluorine) to 0.7 (Cesium).
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Figure 8.3
The Pauling Electronegativity Values
Electronegativity values range from 4.0 (F) to 0.7 (Cs).
Trend??
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EN
Electronegativity Difference
• EN = 0
Non-polar Covalent
identical atoms
• EN = 0.5 to 1.9
• EN > 1.9
Polar Covalent
Ionic
• Approximate Values
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Problem
• Using the Periodic Table, order the following
from lowest to highest electronegativity.
• Fr, Mg, Rb
Fr < Rb < Mg
• P, As, Ga, O
Ga < As < P < O
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Section 8.3
Bond Polarity & Dipole Moments
•
A molecule, such as HF, that has a center of
positive charge and a center of negative charge
is said to be polar, or to have a dipole moment.
Molecule said to be dipolar.
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Diatomic Molecules
• Polarity clear cut:
• Cl ― Cl
non-polar,
no dipole moment
electrons equally shared
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Polyatomic Molecules
• More difficult
• Determine the polarity of each bond.
• Consider the 3-D shape.
• Determine direction (if any) of the molecules dipole
moment.
• Sometimes individual dipoles cancel each other out.
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Figure 8.4
Dipole Moment for H2O
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21
Figure 8.5
Dipole Moment for NH3
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22
Figure 8.6
(a) Carbon Dioxide
b) Opposed Bond Polarities
CO2 is Nonpolar
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23
Section 8.4
Ions: Electronic Configurations
and Sizes
Atoms in Stable compounds have a
Noble Gas Electronic Configuration.
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Achieving Noble Gas Electron
Configurations (NGEC)
•
Two nonmetals react: They share electrons
to achieve NGEC.
•
A nonmetal and a representative group
metal react (form ionic compound): The
valence orbitals of the metal are emptied to
achieve NGEC. The valence electron
configuration of the nonmetal achieves NGEC
by gaining these electrons. .
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Atoms lose or gain electrons to achieve a
NGEC.
Example:
Mg 1s2 2s2 2p6 3s2
lose 2 electrons
Mg2+ 1s2 2s2 2p6
 isoelectric with Ne
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Example:
P: 1s2 2s2 2p6 3s2 3p3
gains 3 e-
P3-: 1s2 2s2 2p6 3s2 3p6
isoelectric with argon
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Isoelectronic Ions
•
Ions containing the the same number of
electrons
• (O2, F, Na+, Mg2+, Al3+)
• O2> F > Na+ > Mg2+ > Al3+
largest
smallest
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• Know Exceptions
• Sn2+
3+
• Bi
• Pb2+
• Tl1+
Sn4+
5+
Bi
Pb4+
Tl3+ (Thallium)
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Formulas for Binary Ionic Cpds.
• Requires compounds to be
Electrically Neutral.
• Sum of the cation charges
MUST EQUAL
Sum of the anion charges
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Size of Ions
• Absolute ion sizes Impossible to
define.
• Mostly Interested in Trends
A  A+ + e- smaller
A + e-  Alarger
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Figure 8.7
Sizes of
Ions
Related
to
Positions
of the
Elements
in the
Periodic
Table
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32
Consider Isoelectric Ions
• O2-
F-
Na+
Mg2+
Al3+
• All have [Ne] EC.
• All have 10 electrons.
• But they have different #’s of
protons.
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Isoelectric Ions
# p+:
O2-
F-
Na+
Mg2+
Al3+
8
9
11
12
13
10 electrons feel greater attraction as the
positive charge of nucleus increases.
THEREFORE, ions become smaller.
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Sizes of Ions
• For isoelectric ions, Size  as Z 
• Ion size increases as you go down a
group.
• Trend Complicated L  R
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35
.
Figure 8.7
Sizes of Ions
Related to
Positions of
the Elements
in the Periodic
Table
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36
Section 8.5:
Formation of Binary Ionic Compounds
• Many separate processes go into the
formation of an ionic solid.
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Lattice Energy
•
The change in energy when separated
gaseous ions are packed together to form an
ionic solid.
•
M+(g) + X(g)  MX(s)
•
Lattice energy is negative (exothermic)
from the point of view of the system.
Dominate Energy Term
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• Reaction is broken into steps in order to
calculate the energy associated with the
process.
• Use steps where the energy of the process
is known.
• Energy is a state function.
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Formation of an Ionic Solid
• 1.
Sublimation of the solid metal
M(s)  M(g)
• 2.
[endothermic]
Ionization of the metal atoms
M(g)  M+(g) + e
• 3.
[endothermic]
Dissociation of the nonmetal
1/2X
2(g)
 X(g)
[endothermic]
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Formation of an Ionic Solid
(continued)
• 4.
Formation of X ions in the gas phase:
• 5.
Formation of the solid MX
X(g) + e  X(g)
[exothermic]
M+(g) + X(g)  MX(s)
[quite exothermic]
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Consider:
K (s) + ½ Cl2 (g)  KCl (s)
Let’s break this into several steps.
• K (s)  K (g)
sublimation
E = + 64 kJ/mole
• K (g)  K+ (g)
ionization
E1 = +419 kJ/mole
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K (s) + ½ Cl2 (g)  KCl (s)
• ½ Cl2 (g)  Cl (g)
Dissociation of chlorine molecules
Energy to break bond = 240 kJ/mole
Energy to break ½ mole = 120 kJ
• Cl (g)  Cl- (g) Electron affinity
EA = - 349 kJ/mole
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K (s) + ½ Cl2 (g)  KCl (s)
• K+ (g) + Cl- (g)  KCl (s)
Lattice energy = -690 kJ/mole
• Net Energy Change
= sum of all the energy changes
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Hf (Energy of formation)
= +64 kJ + 419 kJ + 120 kJ
- 349 kJ - 690 kJ
= - 436 kJ
Lattice energy is the dominate energy
term.
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45
Consider
Li (s) + ½ F2 (g)  LiF (s)
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Figure 8.8
The
Energy
Changes
Involved
in the
Formation
of Solid
Lithium
Fluoride
from its
Elements
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Figure 8.9
The Structure of
solid Lithium Fluoride
Ions pack together such
that
1) maximize the
attraction of oppositely
charged ions &
2) minimize the
repulsion of identically
charged ions.
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Comparison of the Energy Changes
Involved in the Formation of
Solid Sodium Fluoride (NaF) and
Solid Magnesium Oxide (MgO).
All ions involved are isoelectric
with Neon.
See next slide.
Compare all energies.
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49
Figure 8.10
Comparison of
the Energy
Changes
Involved in the
Formation of
Solid Sodium
Fluoride (NaF)
and Solid
Magnesium
Oxide (MgO)
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Lattice Energy = k(Q1Q2 / r )
•Q1, Q2 = charges on the ions
•Lattice energy has negative sign when Q1 and Q2
have opposite signs.
• LE increases as ionic charges increase.
•Higher charges  More energetically stable crystal
•r = shortest distance between centers of the cations
and anions
•LE increases as the distance between the ions
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Section 8.6
• Partial Ionic Character of
Covalent Bonds
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Figure 8.11
Three Possible
Types of Bonds
(a) Covalent bond
between Identical
Atoms.
(b) Polar Covalent
Bond.
(c) Ionic Bond (no
sharing of electrons)
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How do we tell the difference between ionic
bond & a polar covalent bond?
• % ionic character of a bond =
measured dipole moment of X—Y
calculated dipole moment of X+ Y-
X 100
Based on calculations done in the gas phase
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Figure 8.12
The Relationship Between the Ionic Character of a Covalent Bond and the Electronegativity Difference of the Bonded Atoms
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None reach 100% ionic character.
Ones with more than 50% are normally
considered ionic solids.
These calculations are for discrete pairs of
atoms, i.e., gas phase where individual XY
molecules exist.
Existence of ions in solid is favored by the
multiple ion interactions.
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56
Operational Definition of Ionic
Compound
• Any compound that conducts an
electric current when melted will be
classified as ionic.
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57
Section 8.7
The Covalent Chemical Bond: A Model
• Carefully Read.
• Complete Handout.
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Models
•
Models are attempts to explain how
nature operates on the microscopic
level based on experiences in the
macroscopic world.
• Bonding is a model proposed to
explain molecular stability.
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Fundamental Properties of Models

A model does not equal reality.

Models are oversimplifications, and are
therefore often wrong.

Models become more complicated as they age.

We must understand the underlying
assumptions in a model so that we don’t
misuse it.
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60