Please note links to Figures and Tables
from the text will not work. Due to
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Chapter 3 - Simple Bonding Theory
Figure Chapter Opener- Ice cluster showing the structure of water molecules and
intermolecular interactions.
3.0 Intro to Simple Bonding Theory

I am going to introduce Lewis Structures and VSEPR and Polarity Through Two
handouts from my web site.
The treatment is very similar to the text.

Handout with Lewis Structure Rules, VSEPR, and Polarity
Note: Your book uses the term steric number instead of effective electron pairs
Note: the error on page 2: five effective electron pairs is Trigonal {Bi}pyramidal

VSEPR Chart

Models 360, A useful web site to display molecules (Java must be activated on
your computer)

3dChem - Has many more molecules than Models 360 (Java must be activated
on your computer)

Topics to Cover in this Section
o Rules for drawing simple structures
o Resonance and bond order
o Structures with less than an octet
o Structures with more than an octet
o Determining Formal Charge
 Rule versus using the structure
 Choosing between several resonance structure with inequivalent
formal charges
 When violating the octet rule gives better formal charges
o VSEPR = Valence Shell Electron Pair Repulsion
o Basic Structures and Angles
o Effect of Lone Pairs and multiple Bonds
o Electronegativity Effects on Bond Angles
o Polarity of Molecules
 Definition
 Nonpolar molecules
 Polar molecules
 Physical Consequences
 Hydrogen Bonding
3.1 Lewis Structures

Simple way of Drawing Structures developed by G. N. Lewis.

Uses valence electrons distributed in shared pairs = bonds and unshared pairs =
lone pairs.

(Page 52) Water Lewis Structure

(Page 52) CO2 and C2H2 Lewis Structures illustrating double and triple bonds,
respectively.
3.1.1 Resonance

Resonance occur when there is a choice of where to place multiple Bonds.

Resonance lowers the overall energy. True structure is a hybrid of the resonance
structures.

Figure 3.1 CO32-, and example of equivalent resonance.
3.1.2 Expanded shells (more than an octet)

Figure 3.2 Examples of compounds that have more than an octet.
3.1.3 Formal Charge

Definition
(
# of valence
)(
lone pairs
)(
1/2 bond electrons
)
Formal Charge =
electrons
in free atom
-
electrons
on atom
-
shared by
the atom


Equivalent way.
Pick an atom.
Start with the number of valence electrons.
Subtract all the lone pair electrons
Subtract one electron for each bond.

The best structures:
minimize the number of formal charges
and put negative charges on the most electronegative atoms.

Figure 3.3 Resonance Structures and Formal Charges for SCN-

Table 3.1 Experimental Data for SCN-. Note a proton attaches to N.

Figure 3.4 Resonance Structures and Formal Charges for OCN
Table 3.2 Experimental Data for OCN-. Note a proton attaches 97% to N and 3% to O.
(Not the structure predicted from the ion)

Figure 3.5 Resonance Structures and Formal Charges for CNO-, fulminate, an unstable
compound.

Figure 3.6 Examples where the best formal charges violate the octet rule.
3.1.4 Multiple Bonds in Be and B Compounds

Be and B are electropositive. Double bonds to these atoms create unreasonable
formal charges.

Lewis structures are drawn with less than an octet, although BF3 shows some
double bond character.

Figure 3.7 Predicted and actual structures for BeF2, BeCl2, and BF3

(Page 57) Example of a B structure that follows the octet rule. B formal charge =
-1, N formal charge = +1
3.2 Valence Shell Electron Pair Repulsion Theory

Electrons are negative, so electron pairs repel each other and are most stable.

Steric Number = number of Lonepairs + number of atoms around a central atom.

Figure 3.8 Basic VSEPR Shapes for Steric Numbers from 2 through 8. Mainly
know 2 thorough 6.

Figure 3.9 Conversion of a Cube to a Square Antiprism
3.2.1 Lone Pair Repulsion

As a general rule the space taken up around a central atom is in the order:
Lone pair > triple bond > double bond > single bond > 1/2 a lone pair (in NO 2, for
example)

Figure 3.10 Effect of Lone Pairs for Steric Number = 4

Cases with Steric number = 5
o
Figure 3.11 Axial versus Equitorial (Actual) Placement of the Lone Pair in
SF4. Equitorial has more room.
o
Figure 3.11 Possible placements of two lone pairs in ClF3. Equitorial has
more room.
o


Possible Structures for ClF3 (ax = axial, eq = equitorial)
 ax-eq - Especially poor due to 90° lone pair-lone pair interaction.
 ax-ax - Poor because it has 6 90° interactions
 ax-eq - Best - More room and only 4 90° interactions
Steric number = 6,
o
One lone pair can replace any of the six equivalent positions. IF5, for
Example.
o
Two lone pair are 180° apart to prevent 90° lone pair-lone pair
interactions. XeF4, for Example.
Steric Number = 7
o One lone pair may cause a distortion (XeF6), but some structures are
octahedral.
o
(Page 63) The distribution of two Lone Pairs at 180° in XeF5-
Figure 3.13 Summary of the effects of lone pairs for Steric Number from 2 to 6

3.2.2 Multiple Bonds

Multiple Bonds have Effects Similar to Lone Pairs, but Usually not as Large.

Figure 3.14 Effect of Double Bond on the Bond Angles in (CH3)2=CH2

Figure 3.15 Summary of the Effect of Double Bonds for Steric Numbers 2
through 6

Figure 3.16 Some Examples with both Lone Pairs and Multiple Bonds
3.2.3 Electronegativity and Atomic Size Effects

Electronegativity = Power of an Atom in a Compound to Pull Electrons to Itself

Table 3.3 Methods of Calculating Electronegativity

Table 3.4 Table of Electronegativities used in your Text

If one keeps the Central atom fixed and varies the outer atoms.
The more Electronegative Atom will have the Smallest Bond Angles.
This is because the Atom with Greater Eletronegativity will Pull Electrons from
the Central Atom.

Note: This Trend agrees with Size Predictions F > Cl > Br > I

(Page 67) Central Atom Fixed, Outer Atom Varied.

On the other Hand, If the Central Atom in Varied, and the Outer Atom Remains
Fixed
The more Electronegative Central Atom will have the Largest Bond Angle.
This is because the more Electronegative Central Pulls Electrons to the Center.

This trend also follows the size of the Varying Atoms.

(Page 68) Central Atom Varies, Outer Atoms Fixed

Table 3.4 Additional Examples for Considering the Effects of Electronegativities

In the Trend in Electronegativity and Size Oppose each Other Size may Win

(Page 69) Size and Electronegativity are Opposed. Size Prediction Wins Out

Axial Bond Lengths are Generally Longer than Equitorial Bond Lengths for Steric
Number = 5

Figure 3.17 Axial and Equitorial Bond Distances in Structures with Steric Number
=5

The Less Electronegative Substituent Tends to Occupy the Equitorial Position in
Structures of Steric Number = 5

Figure 3.18 Structure with Groups of Differing Electronegativity for Steric Number
=5

Figure 3.19 More Examples Structure with Groups of Differing Electronegativity
for Steric Number = 5

Figure 3.20 Distortions by Electronegative Groups are less than either Lone Pairs
or Multiple Bonds

Figure 3.21 Difficulties for Predicting Structures when F and CF3 substituents are
Present in Compounds with Steric Number = 5
3.2.4 Ligand Close Packing

In Many Structures the Non-Bonded Distance between Outer Atoms Remains
Relatively Constant

Table 3.6 Illustration of nearly Constant F····F Non-Bonded Distances in Several
Compounds

These Nearly Equal Distances are Explained by Assuming the Outer Atoms
(Ligands) are Touching each Other
This is called Ligand Close Packing

Figure 3.22 How Ligand Close Packing, a Known Bond Angle and Trigonometry
may be Used to Calculate Bond Angles.

sin(1/2 F-N-F angle) = [1/2 (F····F Non-Bonded Distance)]/(N-F bond Distance)
3.3 Molecular Polarity

Dipole moment indicates a separation of the center of positive charge and
negative charge in a molecule

Presence of a dipole (polar molecule) is determined from bond vectors form the
less polar to more polar atom in a bond.

See the handout for details

Diatomics the bond dipole = bond vector. More complicated use the vector sum
of bond dipoles.

Table 3.7 Comparison of experimental, Bond Vector Sum and Computer
Determinations of Dipole Moment

Figure 3.23 Effect of Lone pairs on dipole Moments

Figure 3.24 Three molecule where the Bond Vector cancel out

Non Polar Molecule Intermolecular forces are called Dispersion Forces.
Due to Temporary Dipoles caused by Random fluctuations in the electron
distribution.
3.4 Hydrogen Bonding

Hydrogen Bonding is an Especially Strong Localized Dipole-Dipole interaction

Figure 3.23 Hydrogen Bonding and the Melting Points of Hydrides

Figure 3.26 Hydrogen Bonds and the Structure of Ice

Figure 3.27 Hydrogen Bonds in Proteins α helix and β pleated sheet.

