Chapter 10
The Shapes of Molecules
Lecture Notes by K. Marr
(Silberberg 3rd Edition)
10.1 Depicting Molecules and Ions with Lewis
Structures
10.2 Using Lewis Structures and Bond Energies to
Calculate Heats of Reaction
10.3 Valence-Shell Electron-Pair Repulsion (VSEPR)
Theory and Molecular Shape
10.4 Molecular Shape and Molecular Polarity
Lewis Structures…..
1.
2.
Indicate the kind of bonding and which atoms
are bonded in molecules and polyatomic ions
Do NOT indicate the molecular shape or
structure. However….
• VSEPR theory uses Lewis structures to predict 3-D
structure
Guidelines for
Writing Lewis Structures
1.
2.
3.
4.
5.
6.
7.
Decide which atoms are bonded
Count all valence electrons (account for the charge of ions!!)
Place 2 electrons in each bond
Complete the octets of the atoms attached to the central
atom by adding electrons in pairs
Place any remaining electrons on the central atom in pairs
If the central atom does not have an octet, form double
bonds, or if necessary, a triple bond.
Write the Lewis Structures for ClF5, TeF4, CO32-,
CH3COO1-
The Octet Rule is Often Violated
H, Be, B, Al violate the octet rule (< 8 valence electrons)
1.
e.g. BeCl2, BH3, AlCl3
Nonmetals with a valence shell greater than n = 2
(e.g. P, Cl, Br, I, etc.)
2.
»
May violate the octet rule when they are the CENTRAL
atom (e.g. ClF5 )
– How can they do this?
– Why doesn’t Fluorine violate the octet rule?
Lewis Structures for Organic Compounds
1.
2.
3.
4.
Alkanes: CnH2n+2
» Methane, Ethane, Propane, Butane, Pentane, Hexane
– What are isomers?
Alkenes: CnH2n have double bond(s)
» One double bond: Ethene (ethylene), Propene
(propylene)
Alcohols: CnH2n+1OH have hydroxyl group(s)
» methanol, ethanol
Carboxylic Acids: CnH2n+1COOH
have carboxyl group(s)
» Methanoic acid (formic acid), HCOOH
» Ethanoic acid (acetic acid, CH3COOH
Using Formal Charge to Select
the Favored Lewis Structure
Sometimes more than one Lewis Structure is possible for a
compound e.g. sulfuric acid, H2SO4; phosphate ion, PO4-3
Formal Charge

•
•
•
Apparent charge on a bonded atom
An atom “owns” all of its nonbonding electrons and half of its bonding
electrons.
The Lewis Structure with the lowest total formal charge is favored
Formal charge of atom =
[# valence e-] – [# unshared e- + 1/2 # shared e-]
OR
F.C. = [# of valence e-] - [# of unshared + # bonds formed ]
Use of Formal Charge to
Select the Favored Lewis Structure
Formal Charge
•
•
•
Apparent charge on a bonded atom
An atom “owns” all of its nonbonding electrons and half of its
bonding electrons.
The Lewis Structure with the lowest total formal charge is
favored
Formal charge of atom =
[# valence e-] – [# unshared e- + 1/2 # shared e-]
OR
F.C. = [# of valence e-] - [# of unshared + # bonds formed ]
Use of Formal Charge to
Select the Favored Lewis Structure
Use formal charge to determine the correct Lewis
structure for
a) sulfuric acid, H2SO4
b) phosphate ion, PO4-3
Recall:
F.C. = [# valence e-] – [# unshared e- + 1/2 # shared e-]
OR
F.C. = [# of valence e-] - [# of unshared + # bonds formed ]
Formal Charge
Three criteria for choosing the more important structure
1. Smaller formal charges (either positive or negative)
are preferable to larger charges;
2. Avoid like charges (+ + or - - ) on adjacent atoms;
3. A more negative formal charge should exist on an
atom with a larger EN value.
Resonance
When Lewis Structures Fail.....
1.
2.
3.
Write the Lewis Structure for the nitrate ion, NO3» Based on your Lewis structure, what kind of bonding would be
expected ?
Experimental measurements indicate....
» All bond lengths and energies are the same!! (B.O. = 1.33)
The NO3- is a Resonance Hybrid of 3 different Lewis
structures....
» Just as mule is neither a horse or a donkey, none of the 3
structures represent NO3-
Resonance Hybrids
1.
2.
3.
Each resonance structure does not actually exist!!
The actual molecule or ion is a hybrid or average of each
resonance structure
Electron-Pair Delocalization
a) Each bonding electron pair is delocalized or spread over the
entire molecule or ion.
b) Results in identical bonds with extra stability since electron
repulsions reduced
Resonance Structures:
Practice Makes Perfect?
1.
Draw the resonance structures for the nitrite ion,
NO2 and the phosphite ion, PO3-3
2.
How do you know when to use resonance?
How do you know how many resonance structures are
possible?
Draw the Lewis structures for ......
» The oxalate ion, C2O4-2
» Benzene, C6H6
3.
4.
– Benzene has a hexagonal ring structure
Using Bond Energies to
Calculate Heats of Reaction, DHrxn
Lewis structures can be used to calculate DHrxn
 For a reaction to occur….

» Bonds within the reactants must be broken (endothermic)
» Bonds within the reactants must be made (exothermic)
DHrxn = S DHreactant bonds broken + S DHproduct bonds formed

Reactants and products must be in gaseous state!! Why??
Using Bond Energies to Calculate Heats of Rxn
DHrxn = S DHreactant bonds broken + S DHproduct bonds formed
e.g. CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (g) DH0rxn= -818 kJ/mol
Figure 10.3
Using bond energies to calculate DH0comb. of Methane, CH4
BOND BREAKAGE
4BE(C-H)= +1652kJ
2BE(O2)= + 996kJ
DH0(bond breaking) = +2648kJ
BOND FORMATION
Enthalpy,H
2[-BE(C O)]= -1598kJ
4[-BE(O-H)]= -1868kJ
DH0(bond forming) = -3466kJ
DH0rxn= -818 kJ/mol
Examples:
Using Bond Energies to Calculate Heats of Reaction, DHrxn
Use bond energies (see table 9.2, page 340 3rd ed) to
calculate in kJ/mole the
1. Standard heat of formation of water (compare your
answer with Appendix B—they should be the same)
2. Standard heat of combustion of propane, C3H8
(ans. = -2042 kJ/mol)
a) Now use standard heats of formation, DHof, to calculate
the heat of combustion of propane, C3H8 (ans. = - 2043 kJ/mol)
Predicting the Shapes of Molecules:
VSEPR Theory
1.
Valence Shell Electron Pair Repulsion
Theory
» In order to limit electrostatic repulsion,
electron pairs in the orbitals around the
central atom stay as far apart as possible
VSEPR: A balloon analogy for the mutual repulsion of electron groups.
Linear
Figure 10.4
Trigonal Planar
Trigonal
Bipyramidal
Tetrahedral
Octahedral
VSEPR Theory
The Number of Electron Pairs around the Central
Atom Determine Molecular Geometry....
2 bonding pairs  linear (Bond angle = 180o)
3 bonding pairs  planar triangle (Bond angle = 120o)
4 bonding pairs  tetrahedral (Bond angle = 109.5o)
5 bonding pairs  trigonal bipyramidal (Bond angles =
90o and 120o )
6 bonding pairs  octahedral (Bond angle = 90o)
Figure 10.5
Predicting Molecular Geometry
Use Lewis structures and VSEPR Theory to
explain the following molecular geometries....
a) H2O and SnCl2
Are they Bent or V-shaped molecules?
b) BeCl2 and CO2
Bent or linear molecules?
Treat double bonds as if only one pair...Why?
Predicting Molecular
Geometry
Use Lewis structures and VSEPR Theory to predict the
following molecular geometries....
1. BH3
2. NH3
3. ClF3
4. ClF3: T-Shaped and NOT trigonal planar. Why??
a) Nonbonding pairs take up more space than bonding
electrons......why?
b) Therefore, nonbonding pairs need to be separated as much as
possible.
Predicting Molecular Geometry
Use Lewis structures and VSEPR Theory to predict the
following molecular geometries....
1. CH4 and PO43- (Ans. Tetrahedral)
2. XeF4
(Ans. Square planar. Why not tetrahedral?)
3. PCl5
(Ans. Trigonal bipyramidal)
4. BrF5
(Ans. Square pyramidal)
SAMPLE PROBLEM 10.9
PROBLEM:
PLAN:
Predicting Molecular Shapes with More Than
One Central Atom
Determine the shape around each of the central atoms in
acetone, (CH3)2C=O.
Find the shape of one atom at a time after writing the Lewis
structure.
SOLUTION:
tetrahedral
H
H C
H
O
C
H
C H
H
tetrahedral
trigonal planar
O
H
C
H C
C
H
HH
>1200
H
<1200
Predicting Molecular
Shapes with More Than
One Central Atom
The tetrahedral
centers of ethanol.
Figure 10.13
Figure 10.9
Lewis structures and molecular shapes
Molecular Polarity
1.
2.
3.
Influences Chemical and Physical Properties
Polar molecules have higher MP’s and BP’s than
nonpolar molecules.....Why?
• Magnitude of Dipole moment influences MP and BP
e.g. H2O vs H2S
Solubility: Like dissolves Like
•
•
Polar solutes dissolve in polar solvents
Nonpolar solutes dissolve in nonpolar solvents
Nonpolar Molecules
1.
2.
Any Molecule with only nonpolar bonds
e.g. F2 and C8H18
Symmetrical Molecules with Polar Bonds of equal
dipole moment.....
a) CO2 , BCl3, and CCl4
b) PCl5 and SF6
Polar Molecules
1.
Asymmetrical Molecules with Polar Bonds
a) H2O and NH3
b) HCl
2.
Symmetrical Molecules with Polar Bonds of
unequal dipole moment
e.g. CHCl3 and CF2Cl2
Note: CCl2F2  CFC-12  once used in refrigerators Ozone
depletion
Electronegativities of the Elements
Figure 10.14
The orientation of polar molecules
in an electric field.
Electric field OFF
Electric field ON