Chapter 10 The Shapes of Molecules

Lecture Notes by K. Marr (Silberberg 3 rd 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

1.

2.

3.

4.

5.

6.

7.

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 ClF 5 , TeF 4 , CO 3 2 , CH 3 COO 1-

The Octet Rule is Often Violated

1.

2.

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

»

e.g. BeCl 2 , BH 3 , AlCl 3 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. ClF 5 ) How can they do this?

–

Why doesn’t Fluorine violate the octet rule?

1.

2.

3.

4.

Lewis Structures for Organic Compounds

»

Alkanes: C n H 2n+2 Methane, Ethane, Propane, Butane, Pentane, Hexane

– What are isomers? »

Alkenes: C n H 2n (propylene) have double bond(s) One double bond: Ethene (ethylene), Propene

»

Alcohols: C n H 2n+1 OH methanol, ethanol

»

have hydroxyl group(s)

»

Carboxylic Acids: C n H 2n+1 COOH h ave carboxyl group(s) Methanoic acid (formic acid), HCOOH Ethanoic acid (acetic acid, CH 3 COOH

Using Formal Charge to Select the Favored Lewis Structure

 Sometimes more than one Lewis Structure is possible for a compound e.g.

sulfuric acid, H 2 SO 4; phosphate ion, PO 4 -3 Formal Charge

• • Apparent charge on a

bonded

atom An atom “owns” all of its nonbonding electrons and

half

electrons.

of its bonding • 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

bonding electrons.

• of its 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) b) sulfuric acid, H 2 SO 4 phosphate ion, PO 4 -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.

2.

3.

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.

Resonance When Lewis Structures Fail.....

1.

2.

3.

Write the Lewis Structure for the nitrate ion, NO 3 -

» 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 NO 3 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 NO 3 -

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

a) Electron-Pair Delocalization 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.

2.

3.

4.

Draw the resonance structures for the nitrite ion, NO 2 and the phosphite ion, PO 3 -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, C 2 O 4 -2

»

Benzene, C 6 H 6

– Benzene has a hexagonal ring structure

Using Bond Energies to Calculate Heats of Reaction,

D

H rxn

  Lewis structures can be used to calculate D H rxn For a reaction to occur….

» Bonds within the reactants must be broken (endothermic) » Bonds within the reactants must be made (exothermic) D

H rxn =

S D

H reactant bonds broken +

S D

H product bonds formed



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

Using Bond Energies to Calculate Heats of Rxn

D

H rxn =

S D

H reactant bonds broken +

S D

H product bonds formed e.g. CH 4 (g) + 2 O 2 (g)



CO 2 (g) + 2 H 2 O (g)

D

H 0 rxn = -818 kJ/mol

Figure 10.3

Using bond energies to calculate

D

H 0 comb.

of Methane, CH 4 BOND BREAKAGE

4BE(C-H)= +1652kJ 2BE(O 2 )= + 996kJ D

H 0 (bond breaking) = +2648kJ BOND FORMATION

2[-BE(C O)]= -1598kJ 4[-BE(O-H)]= -1868kJ D

H 0 (bond forming) = -3466kJ

D

H 0 rxn = -818 kJ/mol

Examples:

Using Bond Energies to Calculate Heats of Reaction,

D

H rxn Use bond energies (see table 9.2, page 340 3 rd 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, C 3 H 8 (ans. = -2042 kJ/mol)

a) Now use standard heats of formation,

D

H o f , to calculate the heat of combustion of propane, C 3 H 8

(ans. = - 2043 kJ/mol)

Predicting the Shapes of Molecules: VSEPR Theory

1.

»

V

alence

S

hell

E

lectron

P

air

R

epulsion 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 Trigonal Planar Tetrahedral Figure 10.4

Trigonal Bipyramidal Octahedral

VSEPR Theory

The Number of Electron Pairs around the Central Atom Determine Molecular Geometry.... 2 bonding pairs



linear (Bond angle = 180 o ) 3 bonding pairs



planar triangle (Bond angle = 120 o ) 4 bonding pairs



tetrahedral (Bond angle = 109.5

o ) 5 bonding pairs



90 o and 120 o ) trigonal bipyramidal (Bond angles = 6 bonding pairs



octahedral (Bond angle = 90 o )

Figure 10.5

Predicting Molecular Geometry

Use Lewis structures and VSEPR Theory to explain the following molecular geometries....

a) H 2 O and SnCl 2 Are they Bent or V-shaped molecules?

b) BeCl 2 and CO 2 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.

2.

3.

BH 3 NH 3 ClF 3 4.

ClF 3 : 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. CH 4 and PO 4 3 (Ans.

Tetrahedral)

2. XeF 4 (Ans.

Square planar. Why not tetrahedral?)

3. PCl 5 4. BrF 5 (Ans.

Trigonal bipyramidal)

(Ans.

Square pyramidal)

SAMPLE PROBLEM 10.9

Predicting Molecular Shapes with More Than One Central Atom PROBLEM: PLAN: Determine the shape around each of the central atoms in acetone, (CH 3 ) 2 C=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 trigonal planar

tetrahedral

H H C H O C C H H H >120 0 <120 0

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. H 2 O vs H 2 S 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. F 2 and C 8 H 18

Symmetrical Molecules with Polar Bonds of equal

dipole moment.....

a) b) CO 2 , BCl 3 , and CCl 4 PCl 5 and SF 6

Polar Molecules

1.

2.

a) b)

Asymmetrical Molecules with Polar Bonds

H 2 O and NH 3 HCl

Symmetrical Molecules with Polar Bonds of unequal dipole moment

e.g. CHCl 3 and CF 2 Cl 2 Note :

CCl 2 F 2

depletion 

CFC-12

 once used in refrigerators  Ozone

Electronegativities of the Elements

Figure 10.14

The orientation of polar molecules in an electric field.

Electric field OFF

Electric field ON