Chapter 10 bonding theory and molecular structure
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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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Chapter 10 Bonding Theory and Molecular Structure

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Chapter 10Bonding 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

      • 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


  • 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:


  • Molecular Shapes

    • The VSEPR model

      • molecular geometries


  • Molecular Shapes

    • The VSEPR model

      • molecular geometries

Electron pair geometry: tetrahedral

Molecular geometry:

tetrahedral

trigonal pyramidal

bent


  • Molecular Shapes

    • The VSEPR model

      • molecular geometries

NI3SO2

PCl4–NO3–

OF2SO32–

BrCl3PO43–


  • Molecular Shapes

    • Molecular polarity

Molecular polarity  physical and chemical properties

d+d–

bonds:if DX > 0  polar bondA—B

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

BeCl2BF3CH2O

CCl4CHCl3NH3

dipole


  • Molecular Shapes

    • Molecular polarity

PCl3F2

CO32–

CHO2–


  • Valence Bond Theory

    • Covalent bonding and orbital overlap

Bonds are formed using valence electrons and orbitals:

overlap

atomic orbitalsmolecular orbitals (covalent bonds)

e.g.,


  • 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?


  • Valence Bond Theory

    • Hybrid orbitals

      • sp hybrid orbitals

BeH2 facts:

2 equivalent bonds


  • Valence Bond Theory

    • Hybrid orbitals

      • sp2 hybrid orbitals

BH3 facts:

trigonal planar,

3 equivalent bonds


  • Valence Bond Theory

    • Hybrid orbitals

      • sp3 hybrid orbitals

tetrahedral,

4 equivalent bonds

CH4 facts:


  • Valence Bond Theory

    • Hybrid orbitals

      • sp3 hybrid orbitals


  • Valence Bond Theory

    • Hybrid orbitals

      • hybridization involving d orbitals


  • Valence Bond Theory

    • Hybrid orbitals

Summary:

e– pair geometryhybridization

linearsp

trigonal planarsp2

tetrahedralsp3

trigonal bipyramidalsp3d

octahedralsp3d2


  • Valence Bond Theory

    • Hybrid orbitals

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

CCl4BrCl3

BF3SF6

NH3BeCl2

PCl4–XeF4


  • Valence Bond Theory

    • Multiple bonds

      • double bonds

trigonal planar = sp2

all six atoms lie

in same plane

C2H4 facts:


  • Valence Bond Theory

    • Multiple bonds

      • triple bonds

C2H2 facts:

linear = sp


  • 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?


  • 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.


  • Molecular Orbital Theory

Overlap of wave functions:

constructive

overlap

destructive

overlap


  • 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)


  • 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)


  • Molecular Orbital Theory

    • First-row diatomics

He2

He2+

b.o. = 0

b.o. = 0.5


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)


  • Molecular Orbital Theory

    • Second-row diatomics

E


  • Molecular Orbital Theory

    • Second-row diatomics


  • Molecular Orbital Theory

    • Second-row diatomics


benzene

C6H6

6 e– in a cyclic,

planar p system

 aromatic stabilization

all sp2

120º

  • Benzene and Aromatic Compounds

planar

hexagon


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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