Chapter 6
Fundamentals of Chemical Bonding
6.1 Overview of Bonding
6.2 Lewis Structures
6.3 Molecular Shapes:
Tetrahedral Systems
6.4 Other Molecular Shapes
6.5 Properties of Covalent
Bonds
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6.1 Overview of Bonding
Learning objective:
Use the concept of electronegativity to determine the
polarity of a chemical bond
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6.1 Overview of Bonding
 Electrons and nuclei are continually moving.
 But they arrange themselves in ways that optimize the
net attractive forces among the electrons and the
nuclei.
 The net electrical energy can be calculated.
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Schematic illustration of two electrons and two nuclei arranged so that
attractive coulombic interactions (blue lines) are greater than repulsive
coulombic interactions (red lines).
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Chemical Bond Formation
 Electrons and nuclei in a molecule balance all
interactions to give the molecule stability.
 Balance is achieved when the electrons are
concentrated between the nuclei.
 The electrons are shared between the nuclei and
this sharing is called a covalent bond.
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The Hydrogen Molecule
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Bond Length and Bond Energy
 Bond length – the
separation distance
where the molecule is
most stable
 Bond energy – the
amount of stability at
this separation
distance, also known
as the strength of the
bond.
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Example 6 -1
The bond length of molecular fluorine is 142 pm, and the
bond energy is 155 kJ/mol. Draw a figure similar to
Figure 6 – 2 that includes both F2 and H2. Write a
caption for the figure that summarizes the comparison
of these two diatomic molecules.
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Let’s Look at F2
 The fluorine atom has 7 valence electrons (2s22p5)
 By gaining an electron, it will become isoelectronic with
neon (2s22p6)
 If two fluorine atoms come together, they can share
the 8th electron.
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Let’s Look at F2
What happens to the orbitals with nonbonding electrons?
The orbitals are still there!
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Unequal Electron Sharing
 A pure covalent bond occurs only when two identical
atoms are bonded: N2, H2, F2, etc.
 When two dissimilar atoms form a covalent bond, the
electron pair is unequally shared, the bond is called a
polar covalent bond
 Therefore, the electrons are nearer to one of the
atoms, and that atom acquires a partial negative
charge.
 And consequently the other atom has a partial positive
charge.
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Polar Bonds
 The bond is referred to as polar and the molecule
can be called a dipole (having two poles)
 The Greek symbol delta “d” is used to indicate
partial charge
How do we determine which atom has the partial negative
charge and which atom has the partial positive charge?
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Electronegativity is the Answer!
 Electronegativity – the ability to attract bonding
electrons.
 Denoted by the Greek symbol chi, c
 When two atoms have different electronegativities, the
bond between them is polar.
 The bigger the difference in electronegativities, the
more polar the bond.
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Trends in Electronegativity
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Polar Bonds
 Nonmetals are more electronegative than metals.
 In general: the further apart the atoms are on the
periodic table, the larger the difference in
electronegativity.
 And, the larger the difference in electronegativity, the
more polar the bond.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 6 - 2
Use the periodic table, without looking up
electronegativity values, to rank each set of three
bonds from least polar to most polar:
(a) S – Cl, Te – Cl, Se – Cl; and (b) C – S, C – O, and C – F.
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6.2 Lewis Structures
Learning objective:
Draw optimized Lewis structures of covalent compounds,
including resonance structures
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6.2 Lewis Structures
 Convenient representations of valence electrons
 Consists of the chemical symbol for the element plus a
dot for each valence electron.
 In normal circumstances, 2 electrons per side, 4 sides.
 If all sides are full, 8 electrons are in the valence
shell…this is called an octet
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The Conventions
Follow the steps for drawing the Lewis Dot Structure of HF
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The Conventions
Follow the steps for drawing the Lewis Dot Structure of HF
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
The Conventions
Follow the steps for drawing the Lewis Dot Structure of HF
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The Conventions
Follow the steps for drawing the Lewis Dot Structure of HF
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The Bonding Framework
a.
b.
c.
d.
e.
An outer atom bonds to only one other atom. An
inner atom bonds to more than one other atom
Hydrogen atoms are always outer atoms.
In inorganic compounds, outer atoms other than
hydrogen usually are the ones with the highest
electronegativities.
The order in which atoms appear in the formula
often indicates the bonding pattern
The hydrogen atoms appear first in the formula of
oxoacid. Nevertheless, in almost all cases these
acidic hydrogen atoms bond to oxygen atoms, not to
the central atom.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Building the Lewis Structure
1. Count the valence electrons.
2. Assemble the bonding framework, placing two electrons per bond.
3. Complete the octets on each outer atom, except H.
4. Assign the remaining electrons to inner atoms.
5. Optimize electron configurations of the inner atoms.
6. Identify equivalent or near-equivalent Lewis structures.
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e.g. PCl3
5 + (3 x 7) = 26 eBonding Pairs
Lone Pairs (nonbonding electrons)
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Example 6 - 3
Determine the provisional Lewis structure of the BF4anion.
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Example 6 - 4
Determine the provisional Lewis structure of
diethylamine, (CH3CH2)2NH.
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Optimizing the Structure
Step 5: Optimize electron configurations of inner atoms.
Check to see if any inner atoms lacks an octet. If
needed, move electrons from adjacent outer atoms
to make double or triple bonds until the octet is
complete.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 6 - 5
Aqueous solutions of formaldehyde, H2CO, are used to
preserve biological specimens. Determine the Lewis
structure of formaldehyde.
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Example 6 - 6
Acrylonitrile, H2CCHCN, is used to manufacture polymers
for synthetic fibers. Draw the Lewis structure of
acrylonitrile.
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Beyond the Octet
 Elements in the 3rd period or higher can have more
than an octet if needed.
 Atoms of these elements have valence d orbitals, which
allow them to accommodate more than eight electrons.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 6 - 7
Chlorine trifluoride is used to recover uranium from
nuclear fuel rods in a high temperature reaction that
produces gaseous uranium hexafluoride
2 ClF3 (g) + U (s) → UF6 (g) + Cl2 (g)
Determine the Lewis structure of ClF3
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Formal Charge
The difference between the number of valence electrons
in the free atom and the number of electrons assigned
to that atom in the Lewis structure.
FC  (Valence electrons in the free atom) (Valence electrons assigned to that atom in the Lewis structure)
If Step 4 leads to a positive formal charge on an inner
atom beyond the second row, shift electrons to make
double or triple bonds to minimize formal charge, even
if this gives an inner atom with more than an octet of
electrons.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 6 - 8
As described in Chapter 2, sulphur dioxide, a by-product
of burning fossil fuels, is the primary contributor to acid
rain. Determine the Lewis structure of SO2.
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Example 6 - 9
Acetic Acid (CH3CO2H, a carboxylic acid) is an important
industrial chemical and is the sour ingredient in
vinegar. Build its Lewis structure.
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Resonance Structures
 Step 6: Identify equivalent or near-equivalent Lewis
structures
 Let’s look at nitrate, NO3-
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Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 6 - 10
Determine the Lewis structure of dihydrogen phosphate,
H2PO4-.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 6 - 11
Determine the Lewis structure of dinitrogen oxide (NNO),
a gas used as an anaesthetic, a foaming agent, and a
propellant for whipped cream.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Hints on Lewis Dot Structures
1.
2.
3.
4.
5.
6.
7.
Octet rule is the most useful guideline.
Carbon forms 4 bonds.
Hydrogen typically forms one bond to other atoms.
When multiple bonds are forming, they are usually
between C, N, O or S.
Nonmetals can form single, double, and triple bonds,
but not quadruple bonds.
Always account for single bonds and lone pairs
before forming multiple bonds.
Look for resonance structures.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
6.3 Molecular Shapes: Tetrahedral Systems
Learning objective:
Recognize the importance of the tetrahedral shape in
molecules
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6.3 Molecular Shapes: Tetrahedral Systems
 Molecules have three dimensional shapes.
 The 3-D shapes define the properties of the molecules.
 How do we predict the shapes?
 VSEPR Theory – valence shell electron-pair repulsion
theory

Electron pairs in the outer shell of an atom repel one another
and end up as far away from each other as possible.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Let’s Take a Step Back…
 Molecules have 3-D shapes because orbitals have 3-D
shapes.
Let’s look at methane:
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Methane, CH4
 In the plane of the paper, it looks like the bond angles
are 90°, but, we know that the molecule exists in three
dimensions.
 The shape is called tetrahedral and has bond angles of
109.5°.
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Carbon and the Tetrahedron
 Hydrocarbons – molecules that contain only carbon
and hydrogen
 Alkanes – hydrocarbons in which each carbon atom
forms bonds to four other atoms.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
The VSEPR Model
First some definitions:
Electron group – a set of electrons that occupies a particular
region around an atom.
 Ligand – an atom or a group of atoms bonded to an inner atom
 Steric number – the sum of the number of ligands plus the
number of lone pairs; in other words, the total number of
groups associated with that atom.

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Molecular Shape
All molecules above have the same
steric number or electron group geometry
(3-D arrangement of the valence shell electron groups)
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Molecular Shape
The molecular shape describes how the ligands (not
the electron groups) are arranged in space.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Determining Molecular Shape
1. Determine the Lewis structure.
2. Use the Lewis structure to find steric numbers for inner atoms.
3. Determine the electron group geometries from the steric
numbers.
4. Use the ligand count to derive molecular shapes from electron
group geometries.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 6 - 12
Describe the shape of the hydronium ion (H3O+). Make a
sketch of the ion that shows the three-dimensional
shape, including any lone pairs that may be present.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 6 - 13
Describe the shape of hydroxylamine, HONH2.
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Silicon
 Silicon displays tetrahedral
shape to virtually all of its
stable compounds.
 95% of crustal rock and its
various decomposition
products are composed of
silicon oxides.
 The principle oxide of
silicon is silica, with
empirical formula SiO2.
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6.4 Other Molecular Shapes
Learning objective:
Use the VSEPR model to predict the shapes of molecules
with steric numbers 2, 3, 5 and 6
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6.4 Other Molecular Shapes
 Steric Number 2: Linear Electron Group Geometry
 Steric Number 3: Trigonal Planar Electron Group
Geometry
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6.4 Other Molecular Shapes (cont.)
 Steric Number 5: Trigonal Bipyramidal Electron
Group Geometry
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6.4 Other Molecular Shapes (cont.)
 Steric Number 6: Octahedral Electron Group
Geometry
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Example 6 - 15
Describe the geometry and draw a ball-and-stick sketch
of Xenon tetrafluoride.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
6.5 Properties of Covalent Bonds
Learning objective:
Understand the factors that influence bond angles,
lengths and energies
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
6.5 Properties of Covalent Bonds
 Bond angles – each of the steric
groups results in well-defined bond
angles.
 When the steric number of an atom
changes, bond angles change exactly
as the model predicts.
 Lone pairs in a molecule cause bond
angles to be a few degrees smaller
than predicted for symmetrical
geometry.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 6 - 16
Experiments show that
sulphur tetrafluoride
has bond angles of
86.9° and 101.5°.
Give an interpretation of
these bond angles.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Dipole Moments
 Polar bonds can result in polar molecules, depending
on the molecule’s geometry.
 A polar molecule will align itself in an electric field.
 The extent to which the molecules align in a field is
referred to as the dipole moment and has the Greek
symbol mu, m.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 6 - 17
Does either ClF5 or XeF4 have a dipole moment?
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Bond Lengths and Energies
 Two important properties of bonds to study:
 Bond
length – the nuclear separation distance where
the molecule is most stable.
 Bond
energy – the stability of a chemical bond.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Table 6 – 1: Average Bond Lengths
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Table 6 – 1: Average Bond Lengths
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What Affects Bond Length?
1.
2.
3.
4.
The smaller the principle quantum numbers of the
valence orbitals, the shorter the bond.
The higher the bond multiplicity, the shorter the
bond.
The higher the effective nuclear charge of the
bonded atoms, the shorter the bond.
The larger the electronegativity difference, the
shorter the bond.
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Example 6 - 18
What factors account for each of the following
differences in bond length?
a. I2 has a longer bond than Br2.
b. C – N bonds are shorter than C – C bonds.
c. H – C bonds are shorter than C ≡ O
d. The carbon – oxygen bond in formaldehyde, H2C=O,
is longer than the bond in carbon monoxide, C ≡ O.
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Bond Energy
Bond strength increases as more electrons are shared
between the atoms
2. Bond strength increases as the electronegativity
difference (∆χ) between bonded atoms increases.
3. Bond strength decreases as bonds become longer.
1.
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Table 6 – 2 Features of Molecular Geometries
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Table 6 – 2 Features of Molecular Geometries
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Chapter 6 Visual Summary
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Chapter 6 Visual Summary
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Chapter 6 Visual Summary
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Chapter 6 Visual Summary
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.
Chapter 6 Visual Summary
Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.