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Chapter 10 Bonding Theory and Molecular Structure. Molecular Shapes The VSEPR model electron-pair geometries molecular geometries Molecular polarity Valence Bond Theory Covalent bonding and orbital overlap Hybrid orbitals sp hybrid orbitals sp 2 hybrid orbitals sp 3 hybrid orbitals

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slide2
Molecular Shapes
    • The VSEPR model
      • electron-pair geometries
      • molecular geometries
    • Molecular polarity
  • Valence Bond Theory
    • Covalent bonding and orbital overlap
    • Hybrid orbitals
      • sp hybrid orbitals
      • sp2 hybrid orbitals
      • sp3 hybrid orbitals
      • hybridization involving d orbitals
    • Multiple bonds
      • double bonds
      • triple bonds
  • Molecular Orbital Theory
    • First-row diatomics
    • Second-row diatomics
  • Benzene and Aromatic Compounds
slide3
Molecular Shapes
    • The VSEPR model
      • electron-pair geometries

Valence Shell Electron Pair Repulsion Theory: regions of electron density

(single, double, or triple bonds or lone pairs) arrange themselves around an

atom to be as far apart as possible (electron pair repulsion).

Electron pair geometries:

slide5
Molecular Shapes
    • The VSEPR model
      • molecular geometries
slide6
Molecular Shapes
    • The VSEPR model
      • molecular geometries

Electron pair geometry: tetrahedral

Molecular geometry:

tetrahedral

trigonal pyramidal

bent

slide11
Molecular Shapes
    • The VSEPR model
      • molecular geometries

NI3 SO2

PCl4– NO3–

OF2 SO32–

BrCl3 PO43–

slide12
Molecular Shapes
    • Molecular polarity

Molecular polarity  physical and chemical properties

d+d–

bonds: if DX > 0  polar bond A—B

molecules and ions: if dipoles do not exactly cancel, molecule will be polar

BeCl2 BF3 CH2O

CCl4 CHCl3 NH3

dipole

slide13
Molecular Shapes
    • Molecular polarity

PCl3F2

CO32–

CHO2–

slide14
Valence Bond Theory
    • Covalent bonding and orbital overlap

Bonds are formed using valence electrons and orbitals:

overlap

atomic orbitals molecular orbitals (covalent bonds)

e.g.,

slide15
Valence Bond Theory
    • Covalent bonding and orbital overlap

But what about CH4?

Tetrahedral, all bonds

equivalent. How do we

get this from s and p a.o.s?

slide16
Valence Bond Theory
    • Hybrid orbitals
      • sp hybrid orbitals

BeH2 facts:

2 equivalent bonds

slide17
Valence Bond Theory
    • Hybrid orbitals
      • sp2 hybrid orbitals

BH3 facts:

trigonal planar,

3 equivalent bonds

slide18
Valence Bond Theory
    • Hybrid orbitals
      • sp3 hybrid orbitals

tetrahedral,

4 equivalent bonds

CH4 facts:

slide19
Valence Bond Theory
    • Hybrid orbitals
      • sp3 hybrid orbitals
slide20
Valence Bond Theory
    • Hybrid orbitals
      • hybridization involving d orbitals
slide21
Valence Bond Theory
    • Hybrid orbitals

Summary:

e– pair geometry hybridization

linear sp

trigonal planar sp2

tetrahedral sp3

trigonal bipyramidal sp3d

octahedral sp3d2

slide22
Valence Bond Theory
    • Hybrid orbitals

What is the hybridization of the central atom in each of the following?

CCl4 BrCl3

BF3 SF6

NH3 BeCl2

PCl4– XeF4

slide23
Valence Bond Theory
    • Multiple bonds
      • double bonds

trigonal planar = sp2

all six atoms lie

in same plane

C2H4 facts:

slide24
Valence Bond Theory
    • Multiple bonds
      • triple bonds

C2H2 facts:

linear = sp

slide25
Valence Bond Theory

What is the hybridization of each indicated atom in the following

molecule? How many sigma and pi bonds are in the molecule?

slide26
Molecular Orbital Theory

Fact: O2 is paramagnetic!

Lewis structure

VSEPR

Valence bond theory

  • sp2 hybridized
  • lone pairs in sp2 hybrid orbitals
  • bonding pairs in s and p bonds

All show

all electrons

paired.

slide27
Molecular Orbital Theory

Overlap of wave functions:

constructive

overlap

destructive

overlap

slide28
Molecular Orbital Theory
    • First-row diatomics

Overlap of 1s orbitals:

s*1s

antibonding m.o.

(higher energy than

separate atoms)

s1s

bonding m.o.

(lower energy than

separate atoms)

slide29
Molecular Orbital Theory
    • First-row diatomics

(no. of e– in bonding m.o.s) - (no. of e– in antibonding m.o.s)

2

bond order =

H2

b.o. = 1 (i.e., lower energy than separate atoms)

slide30
Molecular Orbital Theory
    • First-row diatomics

He2

He2+

b.o. = 0

b.o. = 0.5

slide31
z

z

x

x

y

y

  • Molecular Orbital Theory
    • Second-row diatomics

Overlap of 2s and 2p orbitals

2s s2s and s*2s

(same as 1s),

then 2p orbitals give:

(i.e., 8 a.o.s  8 m.o.s)

slide32
Molecular Orbital Theory
    • Second-row diatomics

E

slide33
Molecular Orbital Theory
    • Second-row diatomics
slide34
Molecular Orbital Theory
    • Second-row diatomics
slide35
benzene

C6H6

6 e– in a cyclic,

planar p system

 aromatic stabilization

all sp2

120º

  • Benzene and Aromatic Compounds

planar

hexagon

slide36
naphthalene

benzo[a]pyrene

(carcinogen)

p-dichlorobenzene

  • Benzene and Aromatic Compounds

methylbenzene

toluene

1,2-dimethylbenzene

ortho-dimethylbenzene

(o-xylene)

(meta-xylene)

(para-xylene)

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