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. Indicate the kind of bonding and which atoms are bonded in molecules and polyatomic ions

2. 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. Decide which atoms are bonded2. Count all valence electrons (account for the charge of ions!!)

3. Place 2 electrons in each bond4. Complete the octets of the atoms attached to the central

atom by adding electrons in pairs5. Place any remaining electrons on the central atom in

pairs6. If the central atom does not have an octet, form double

bonds, or if necessary, a triple bond.7. Write the Lewis Structures for ClF5, TeF4, CO3

2-, CH3COO1-

The Octet Rule is Often Violated

1. H, Be, B, Al violate the octet rule (< 8 valence electrons)

e.g. BeCl2, BH3, AlCl3

2. 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?

Lewis Structures for Organic Compounds

1. Alkanes: CnH2n+2

» Methane, Ethane, Propane, Butane, Pentane, Hexane– What are isomers?

2. Alkenes: CnH2n have double bond(s)

» One double bond: Ethene (ethylene), Propene (propylene)

3. Alcohols: CnH2n+1OH have hydroxyl group(s)

» methanol, ethanol4. 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.

ResonanceWhen Lewis Structures Fail.....

1. Write the Lewis Structure for the nitrate ion, NO3-

» Based on your Lewis structure, what kind of bonding would be expected ?

2. Experimental measurements indicate....

» All bond lengths and energies are the same!! (B.O. = 1.33)

3. 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. Each resonance structure does not actually exist!!

2. The actual molecule or ion is a hybrid or average of each resonance structure

3. Electron-Pair Delocalizationa) 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?3. How do you know how many resonance structures are

possible?4. Draw the Lewis structures for ......

» The oxalate ion, C2O4-2

» Benzene, C6H6

– Benzene has a hexagonal ring structure

Using Bond Energies to Calculate Heats of Reaction, Hrxn

Lewis structures can be used to calculate Hrxn

For a reaction to occur….» Bonds within the reactants must be broken (endothermic)» Bonds within the reactants must be made (exothermic)

Hrxn = Hreactant bonds broken + Hproduct bonds formed

Reactants and products must be in gaseous state!! Why??

Using Bond Energies to Calculate Heats of Rxn

Hrxn = Hreactant bonds broken + Hproduct bonds formed

e.g. CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g) H0rxn= -818 kJ/mol

Figure 10.3 Using bond energies to calculate H0comb. of Methane, CH4

Ent

halp

y,H

BOND BREAKAGE

4BE(C-H)= +1652kJ

2BE(O2)= + 996kJ

H0(bond breaking) = +2648kJBOND FORMATION

4[-BE(O-H)]= -1868kJ

H0(bond forming) = -3466kJ

H0rxn= -818 kJ/mol

2[-BE(C O)]= -1598kJ

Examples: Using Bond Energies to Calculate Heats of

Reaction, Hrxn

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, Hof, 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.

Figure 10.4

Linear Trigonal Planar Tetrahedral

Trigonal Bipyramidal

Octahedral

VSEPR TheoryThe 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 Predicting Molecular Shapes with More Than One Central Atom

SOLUTION:

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.

C C C

OH

H

H

HH

H

tetrahedral tetrahedral

trigonal planar

C

O

HC

HHH

CH

H>1200

<1200

Figure 10.13

The tetrahedral centers of ethanol.

Predicting Molecular Shapes with More Than One Central Atom

Figure 10.9 Lewis structures and molecular shapes

Molecular Polarity

1. Influences Chemical and Physical Properties

2. 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

3. Solubility: Like dissolves Like• Polar solutes dissolve in polar solvents

• Nonpolar solutes dissolve in nonpolar solvents

Nonpolar Molecules

1. Any Molecule with only nonpolar bonds

e.g. F2 and C8H18

2. 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 Bondsa) 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