11/7/2014
Ch. 4 - Chemical Bonding:
Understanding Climate Change
Chapter Outline
 4.1 Types of Chemical Bonds
 4.2 Naming Compounds and Writing
Formulas
 4.3 Lewis Structures
 4.4 Electronegativity, Unequal Sharing,
and Polar Bonds
 4.5 Vibrating Bonds and the Greenhouse Effect
 4.6 Resonance
 4.7 Formal Charge: Choosing among Lewis Structures
 4.8 Exceptions to the Octet Rule
 4.8 The Lengths and Strengths of Covalent Bonds
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Types of Chemical Bonds
Ionic
Covalent
Metallic
Chemical Bonds

Types of chemical bonds:
•
•
•
Ionic Bond – electrons transferred: chemical
bond resulting from the electrostatic
attraction of a cation for an anion.
Covalent Bond – chemical bond that results
from a sharing of outermost electrons.
Metallic Bond – chemical bond consisting of
the nuclei of metal atoms surrounded by a
“sea” of shared electrons.
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Ionic Bonds
𝑄1 𝑥𝑄2
𝐸𝑒𝑙 ∝
𝑑
Sample Exercise 4.1: Calculating the
electrostatic Potential Energy of an Ionic Bond
What is the electrostatic potential energy (Eel) of the
ionic bond between a potassium ion and a chloride ion?
𝐸𝑒𝑙 = 2.31𝑥10−19 𝐽 ∙ 𝑛𝑚
𝑄1 𝑥𝑄2
𝑑
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Lattice Energy (U)
• The energy released when
one mole of an ionic
compound forms from its free
ions in the gas phase.
𝑄1 𝑥𝑄2
𝐸𝑒𝑙 ∝
𝑑
• As charges increase, so does the
lattice energy
• As the distance between ions
decreases, the lattice energy
increases
Covalent Bonds
H2
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Metallic Bonds
Cu = [Ar]3d104s1
 

3d10



4s1
Bond Type Summary
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Chapter Outline
 4.1 Types of Chemical Bonds
 4.2 Naming Compounds and Writing
Formulas (Lab)
 4.3 Lewis Structures
 4.4 Electronegativity, Unequal Sharing,
and Polar Bonds
 4.5 Vibrating Bonds and the Greenhouse Effect
 4.6 Resonance
 4.7 Formal Charge: Choosing among Lewis Structures
 4.8 Exceptions to the Octet Rule
 4.8 The Lengths and Strengths of Covalent Bonds
1
1
Sec. 4.3: Lewis Bonding Theory:
Lewis Symbols for Atoms
 Valence electrons are drawn as dots around
the symbol
••
•• •
•
 Element symbol = nucleus + core electrons
O
 Up to 4 valence electrons are placed around
the symbol one at a time; additional electrons
are paired up
 The result is up to 4 pairs of electrons = octet
H
•
 NOTE: hydrogen can not have an octet. When
forming bonds with other atoms, it can have a
maximum of 2 electrons in its valence shell
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Lewis Symbols and the Periodic Table
Group
e- configuration
# of valence e-
1A
ns1
1
2A
ns2
2
3A
ns2np1
3
4A
ns2np2
4
5A
ns2np3
5
6A
ns2np4
6
7A
ns2np5
7
Lewis Dot Symbol
Lewis Symbols and the Periodic Table
Unpaired dots = bonding
capacity.
Main Group Elements:
Members of same family
have same number of
valence electrons, and
similar bonding capacities.
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Lewis Structures of Ionic Compounds
Na
[Ne]3s1
e- + Cl
[Ne]2s22p5
Na+ + e[Ne]
Cl
-
Na+ + Cl
-
Na+ Cl
-
[Ne]3s23p6 = [Ar]
Sample Exercise 4.8 (Modified)
Draw the Lewis symbols of the monatomic ions
formed by calcium and oxygen. Then draw the
Lewis structure of calcium oxide (CaO).
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Lewis Structures of Molecular
(Covalent) Compounds
A covalent bond is a chemical bond in which two or
more electrons are shared by two nonmetals,
resulting in an octet for both atoms. For example Lewis structure of F2
Lewis structure of H2O
Guidelines for Drawing Lewis Structures
(updated later on with the concept of “formal charge”)
1. Hydrogen is always a terminal atom because it
can form only one bond.
2. The CENTRAL ATOM usually has the lowest
electron affinity (or electronegativity as defined
later)
3. Arrange the atoms geometrically and
symmetrically.
e.g. CHCl3
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Guidelines for Drawing Lewis Structures
(updated later on with the concept of “formal charge”)
4. Sum up the total number of valence electrons
(use the group number), and calculate the
number of pairs.
5. Connect the atoms together so that each atom
has an octet (except H). You may have to form
multiple bonds.
Multiple Bonds – sharing more than one
pair of electrons
Double bond – two atoms share two pairs of electrons
CO2
H2CO
Triple bond – two atoms share three pairs of electrons
N2
C2H2
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Lewis Structures of Charged Species
ClO-
NO2+
Electronegativity, Unequal Sharing, and Polar Bonds
 Electronegativity ():
• Ability of an atom to attract bonding electrons.
• Periodic trend similar to ionization energy.
Electronegativities
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Ionization Energies
and Electronegativies
EN increases
across a row.
EN decreases
down a column.
Polar Covalent Bonds
• Unequal sharing of electrons in a covalent bond
resulting in an uneven distribution of charge.
• Results from differences in electronegativity.
• Dipole Moment = polarity indicated by arrow
pointing to more “partially negative” end, with a
“partially positive” charge on the opposite size
𝛿+
𝛿−
H
Cl
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Polar Covalent Bonds - Unequal sharing of electrons
resulting in an uneven distribution of charge.
+1
𝛿+
𝛿−
-1
Difference
0 < 0.4
Cl2
Bond Type
Covalent
HCl
0.4 < EN < 2
Polar Covalent
NaCl
2
Ionic
Electronegativity Trends
 As seen previously, electronegativity increases moving
up to the right in the periodic table. (Noble gases not
included.)
 Bond polarity increases as ΔEN increases.
ΔEN = 1.9
0.9
0.7
0.4
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Sample Exercise 4.12
Rank, in order of increasing polarity, the bonds
formed between O and C
Cl and Ca
N and S
O and Si
Are any of these bonds considered ionic?
Chapter Outline









4.1 Types of Chemical Bonds
4.2 Naming Compounds and Writing Formulas
4.3 Lewis Structures
4.4 Electronegativity, Unequal Sharing, and Polar Bonds
4.5 Vibrating Bonds and the Greenhouse Effect
4.6 Resonance
4.7 Formal Charge: Choosing among Lewis Structures
4.8 Exceptions to the Octet Rule
4.8 The Lengths and Strengths of Covalent Bonds
2
8
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The Concept of “Resonance”
Resonance structures (or hybrids) = some molecules have
two or more plausible Lewis structures. The actual structure is
“intermediate” or an “average” of all the possibilities.
e.g. O3
Resonance Structures: Ozone
 There are two plausible
Lewis structures for
ozone.
 The actual structure for O3
is the “average” of the two
resonance structures.
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Sample Exercise 4.13: Drawing
Resonance Structures
Draw all the resonance structures of the nitrate ion.
Resonance in Organic Compounds
Benzene = C6H6 and it forms a ring
3
2
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Chapter Outline









4.1 Types of Chemical Bonds
4.2 Naming Compounds and Writing Formulas
4.3 Lewis Structures
4.4 Electronegativity, Unequal Sharing, and Polar Bonds
4.5 Vibrating Bonds and the Greenhouse Effect
4.6 Resonance
4.7 Formal Charge: Choosing among Lewis Structures
4.8 Exceptions to the Octet Rule
4.8 The Lengths and Strengths of Covalent Bonds
3
3
Formal Charge and Lewis Structures
 Three possible Lewis structures. Which
one is best?
 Formal charge:
• Determined by the difference between the
number of valence e- in the free atom, and
the sum of lone pair + ½ bonding e- in a
molecule.
3
4
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Calculating Formal Charges
 Formal charge (FC):
FC = (# valence electrons) –
[ (# unshared e-) + 1/2(# of e- in bonding pairs)]
3
5
Choosing the Best Lewis Structure
 Most Stable Resonance Structures:
• Formal charges equal or close to zero.
• Negative formal charges on the more
electronegative element.
3
6
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Practice: Formal Charge
What is the most likely Lewis structure for CO2?
Summary: Formal Charge and Lewis Structures
A way to decide which Lewis Structure, out of several
possibilities, that’s the most stable and therefore the most likely.
1. For neutral molecules, a Lewis structure in which there
are no formal charges is preferable to one in which
formal charges are present.
2. Lewis structures with large formal charges are less
plausible than those with small formal charges.
3. Among Lewis structures having similar distributions of
formal charges, the most plausible structure is the one in
which negative formal charges are placed on the more
electronegative atoms.
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Chapter Outline









4.1 Types of Chemical Bonds
4.2 Naming Compounds and Writing Formulas
4.3 Lewis Structures
4.4 Electronegativity, Unequal Sharing, and Polar Bonds
4.5 Vibrating Bonds and the Greenhouse Effect
4.6 Resonance
4.7 Formal Charge: Choosing among Lewis Structures
4.8 Exceptions to the Octet Rule
4.8 The Lengths and Strengths of Covalent Bonds
3
9
Exceptions to the Octet Rule
1. The Incomplete Octet (Central atom in Group 3A)
e.g. BF3
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Exceptions to the Octet Rule
2. Odd-Electron Molecules – “free radicals” (one
unpaired electron = high reactivity)
NO
NO2
Exceptions to the Octet Rule
3. Expanded Octet – central atom has greater than
8 valence electrons surrounding it. Occurs only with
elements in row 3 and higher because they have
available d-orbitals
A. Molecular (Covalent) Molecules
SF6
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Exceptions to the Octet Rule
B. Polyatomic Ions
PO43-
SO42-
Chapter Outline









4.1 Types of Chemical Bonds
4.2 Naming Compounds and Writing Formulas
4.3 Lewis Structures
4.4 Electronegativity, Unequal Sharing, and Polar Bonds
4.5 Vibrating Bonds and the Greenhouse Effect
4.6 Resonance
4.7 Formal Charge: Choosing among Lewis Structures
4.8 Exceptions to the Octet Rule
4.8 The Lengths and Strengths of Covalent Bonds
4
4
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The Lengths and Strengths of Covalent Bonds:
Bond Length vs Bond Order
Bond
Bond
Order
Bond
Length
(pm)
Bond
Energy
(kJ/mol)
C-C
1
154
348
C=C
2
134
614
CC
3
120
839
C-N
1
143
293
C=C
2
138
615
CN
3
116
891
Resonance Structures have an
intermediate bond order
Bond order = 1.5
Bond order = 2
Bond order = 1
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Table of Average Covalent Bond Lengths and
Bond Energies – can be used to estimate Hrxn
Bond energy = the average energy required to
break a particular type of bond.
Bond Energy
H2 (g)
H (g) + H (g)
H0 = 436 kJ
Cl2 (g)
Cl (g) + Cl (g)
H0 = 243 kJ
HCl (g)
H (g) + Cl (g)
H0 = 431 kJ
O2 (g)
O (g) + O (g)
H0 = 495 kJ
O
O
N2 (g)
N (g) + N (g)
H0 = 941 kJ
N
N
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Make = exothermic
Break = endothermic
H0 = total energy input – total energy released
= SBE(reactants) – SBE(products)
e.g. CH4 + 2 O2  H2O + 2 CO2
Use bond energies to calculate the enthalpy change for:
CH4 + 2 O2  H2O + 2 CO2
H0 =
SBE(reactants) – SBE(products)
Type of
bonds broken
Number of
bonds broken
Bond energy
(kJ/mol)
Energy
change (kJ)
Type of
bonds formed
Number of
bonds formed
Bond energy
(kJ/mol)
Energy
change (kJ)
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