1. Chapter 10 The Shapes of Molecules Lecture Notes by K
Chapter 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

2. Lewis Structures…..
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

3. Guidelines for Writing Lewis Structures
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-

5. The Octet Rule is Often Violated
H, Be, B, Al violate the octet rule (< 8 valence electrons)
e.g. BeCl2, BH3, AlCl3
Nonmetals with a valence shell greater than n = 2 (e.g. P, Cl, Br, I, etc.)
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?

6. Lewis Structures for Organic Compounds
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

7. 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 ]

8. Use of Formal Charge to Select the Favored Lewis Structure
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 ]

9. Use of Formal Charge to Select the Favored Lewis Structure
Use formal charge to determine the correct Lewis structure for
sulfuric acid, H2SO4
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 ]

10. Formal Charge Three criteria for choosing the more important structure
Smaller formal charges (either positive or negative) are preferable to larger charges;
Avoid like charges (+ + or - - ) on adjacent atoms;
A more negative formal charge should exist on an atom with a larger EN value.

11. Resonance When Lewis Structures Fail.....
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-

13. Resonance Hybrids Each resonance structure does not actually exist!!
The actual molecule or ion is a hybrid or average of each resonance structure
Electron-Pair Delocalization
Each bonding electron pair is delocalized or spread over the entire molecule or ion.
Results in identical bonds with extra stability since electron repulsions reduced

14. Resonance Structures: Practice Makes Perfect?
Draw the resonance structures for the nitrite ion, NO2 - and the phosphite ion, PO3-3
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
Benzene has a hexagonal ring structure

15. 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??

16. 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

17. Using bond energies to calculate DH0comb. of Methane, CH4
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
2[-BE(C O)]= -1598kJ
Enthalpy,H
4[-BE(O-H)]= -1868kJ
DH0(bond forming) = -3466kJ
DH0rxn= -818 kJ/mol

18. 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
Standard heat of formation of water (compare your answer with Appendix B—they should be the same)
Standard heat of combustion of propane, C3H (ans. = kJ/mol)
Now use standard heats of formation, DHof, to calculate the heat of combustion of propane, C3H8 (ans. = kJ/mol)

20. Predicting the Shapes of Molecules: VSEPR Theory
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

21. Linear Trigonal Planar Tetrahedral Trigonal Bipyramidal Octahedral
VSEPR: A balloon analogy for the mutual repulsion of electron groups.
Linear
Trigonal Planar
Tetrahedral
Trigonal Bipyramidal
Octahedral
Figure 10.4

22. 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)

23. Figure 10.5

24. Predicting Molecular Geometry
Use Lewis structures and VSEPR Theory to explain the following molecular geometries....
H2O and SnCl2
Are they Bent or V-shaped molecules?
BeCl2 and CO2
Bent or linear molecules?
Treat double bonds as if only one pair...Why?

26. Predicting Molecular Geometry
Use Lewis structures and VSEPR Theory to predict the following molecular geometries....
BH3
NH3
ClF3
ClF3: T-Shaped and NOT trigonal planar. Why??
Nonbonding pairs take up more space than bonding electrons......why?
Therefore, nonbonding pairs need to be separated as much as possible.

27. Predicting Molecular Geometry
Use Lewis structures and VSEPR Theory to predict the following molecular geometries....
CH4 and PO (Ans. Tetrahedral)
XeF (Ans. Square planar. Why not tetrahedral?)
PCl (Ans. Trigonal bipyramidal)
BrF (Ans. Square pyramidal)

28. SAMPLE PROBLEM 10.9
Predicting Molecular Shapes with More Than One Central Atom
PROBLEM:
Determine the shape around each of the central atoms in acetone, (CH3)2C=O.
PLAN:
Find the shape of one atom at a time after writing the Lewis structure.
SOLUTION:
tetrahedral
tetrahedral
trigonal planar
>1200
<1200

29. The tetrahedral centers of ethanol.
Predicting Molecular Shapes with More Than One Central Atom
The tetrahedral centers of ethanol.
Figure 10.13

38. Lewis structures and molecular shapes
Figure 10.9
Lewis structures and molecular shapes

50. Molecular Polarity 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

51. Nonpolar Molecules Any Molecule with only nonpolar bonds
e.g. F2 and C8H18
Symmetrical Molecules with Polar Bonds of equal dipole moment.....
CO2 , BCl3, and CCl4
PCl5 and SF6

52. Polar Molecules Asymmetrical Molecules with Polar Bonds
H2O and NH3
HCl
Symmetrical Molecules with Polar Bonds of unequal dipole moment
e.g. CHCl3 and CF2Cl2
Note: CCl2F2  CFC-12  once used in refrigerators Ozone depletion

53. Electronegativities of the Elements

54. The orientation of polar molecules in an electric field.
Figure 10.14
The orientation of polar molecules in an electric field.
Electric field ON
Electric field OFF