Two Bonding Theories

• Valence Bond Theory (localized electron model)– Electrons in a molecule still occupy

orbitals of individual atoms.

• Molecular Orbital Theory– Says that atomic orbitals no longer

exist. Molecular orbitals are available for occupation by electrons.

Valence Bond Theory

• Electrons in a molecule still occupy orbitals of individual ATOMS.

• Half-filled orbitals of two atoms overlap.

VB Theory: Formation of H2S

VB Theory: Formation of H2S

VB Theory: Carbon

2s 2p

Would suggest only two bonds possible

VB Theory: Carbon

2s 2p

What if we just bump an electron up?

Would suggest two different types of bonds form (one with an s orbital and three with p orbitals)

Valence Bond Theory• Atomic Orbitals Hybridize.

• These are still atomic orbitals.• Two or more nonequivalent orbitals

of same atom combine.• When X number of atomic orbitals

hybridize, X number of hybrid orbitals will result.

2s 2p

Four atomic orbitals (one s and three p) hybridize into four degenerate hybrid orbitals, as a group called sp3

2sp3

VB Theory: Carbon

How will four degenerate orbitals exist around the central atom?

How will four degenerate orbitals exist around the central atom?

VB Theory: NH3

VB Theory: sp3 hybridization

VB Theory: sp2 hybridization

VB Theory: sp2 hybridization

VB Theory: sp2 hybridization

C2H4, sp2 hybridization

Valence Bond TheoryHybridization is directly connected to

the Electron-pair Geometry

sp3 sp2 sp = linear

sp3d = trig bipryramidalsp3d2 = octahedral

= tetrahedral= trigonal planar

Molecular Orbital Theory

• Says that atomic orbitals no longer exist.

• Molecular orbitals are available for occupation by electrons.

• Atomic orbitals from all atoms overlap and become an equivalent number of molecular orbitals.

Molecular Orbital TheoryConstructive vs. Destructive

Interference

Molecular Orbital Theory

Molecular Orbital Energy Diagram

for H2

Molecular Orbital Energy Diagramfor He2

Molecular Orbital Theory

Molecular Orbital Theory

Molecular Orbital Theory

Molecular Orbital Theory

• Sigma Bond – electron density along line joining two nuclei

- from the side, would look like an s orbital

- first bond that forms between two nuclei (single bond)

bonding: s1s antibonding: s*1s

s*2ps2p

Molecular Orbital Theory

• Pi Bond – formed by sideways overlapping of orbitals, electron density above and below plane of nuclei

– second / third bond that forms between two nuclei (double / triple bonds)

- less overlap means weaker than sigma bonds

sp

C Atom

H

C

HHH

2s 2p

Promotion Step

C Atom

2s 2p

Hybridization step ( leads to energy release after bonding, since more bonds can be formed)

sp3

For 4 groups: tetrahedral

electron pair geometry

C C

Hybridization of Carbon

For 3 groups: trigonal planar

electron pair geometry

H

H

H

H

For 2 groups: linear

electron pair geometry

sp2 pz

C

py

leftover

H

H

sp py, pzTwo hybridized

AO'sleftover

sp

Example: CH4

Use to form 3 single () bonds

Use pz to form 1

pi () bond

H C

Example: C2H4

bond

C

Each C forms three bonds and

one bond. Bond angles are 120o.

H

xy plane

Use to form 2 single ()

bonds

Use py and pz to

form 2 pi() bonds

Example: C2H2

Each C forms two bonds

and two bonds (which are

perpendicular to each other).

Bond angles are 180o.

( requires energy)

Three hybridized AO's

Four hybridized AO's

pzxy plane

C

H

pz

H

C CH H

Use to form 4 single (sigma = ) bonds

sp2

sp2sp2

sp3 sp3

sp3

sp3

C forms four bonds.

Bond angles are 109.5o.

from Susan Piepho,

Sweet Briar College

http://lhs2.lps.org/staff/squiring/chemistry/Chap15_16.2/CarbonHybridization.html

http://lhs2.lps.org/staff/squiring/chemistry/Chap15_16.2/CarbonHybridization.html