Atomic Orbital Hybridization Lewis structure

Valence Bond Theory

Valence Bond Theory

Application: Can be used to explain the bonding of molecules    

Limitation: Doesn't describe geometry 11. 3 –the Hybridization of molecules

Atom must go from ground state"Basic" to excited state by absorbing energy and "Basic" promoting an electron Valence Bond Theory Bond T The number of hybrid orbitals equals the total number of atomic orbitals thatValence are combined. Application: Limitation: Hybridization occurs only when the bonds are being formed beofused toisexplain the Doesn't describe the ge Energy and the Can number orbitals conserved

bonding of molecules

of molecules

Hybridization: process of replacing pure atomic orbitals with reformulated atomic Atomic Orbital sp hybridization orbitals for bonded atoms

Hybridization

Hydrid Orbital: the new orbital Hybrid Orbitals Geometry

Example: BeCl2

Geometric Orientation Linear (V.B. Theory) Trigonal-planar Tetrahedral Atomic Orbital Trigonal-bipyramidal Hybridization Octahedral

Bonding

sp sp2 sp3 sp3d sp3d2

(VSEPR)

sp hybridizati

Example: BeCl

Geometry

Bonding

(VSEPR)

(V.B. Theory)

Geometry (VSEPR) Lewis structure

Example #1: sp hybridization BeCl2

.. Cl ..

Be

.. Cl ..

1. Draw Lewis Dot Diagram

1800

Geometry (VSE Lewis structure 2. Fill in s and p orbitals

.. Cl ..

Be

Linear molecule Cl 3s

.. Cl ..

.. Cl ..

Be

180 0

.. Cl ..

Cl 3s

3. Promote electron into the p orbital

Hybridization predicts two identical bonds, as expected!

2 Linear molecule

Cl 3s

Cl

3p Be sp

Cl 3s

3p

2p

11. 3 – Hybridization

Results in a two sp with p overlaps

Be 2p

2s

2p

Energy

Hybridization C

Be sp (2)

3 hybrid sp2 orbitals

2p

s

s

C

p

s

p

s

p

p (2)

+

sp2

sp2 sp2 (3)

Example #2: BF3 (Trigonal planar molecule) F

ess!

Success!

F 1s

2s

2p

1s

B

Hybridization predicts two identical bonds, as expected!

2p

B sp2

Results in three hybrid sp2 atomic orbital

2s

sp2

p

+ 2 more… ..

3 equivalent

4 equivalent

p

bonds formed at 120o

bonds

H H F

p

C

C

H

F

sp2 sp2

sp2

F

F

F

p

B

B

H

120

0

F

p

2 hybrid sp orbitals 3 equivalent

s

bonds formed at 120o

p

9

One pure s orbital

Example #3: CH4 (Tetrahedral)

p

Three pure p orbitals

Energy + C s

p

Hybridization C

4

= C Four hybrid sp3 atomic orbitals s

p

sp3 (4)

Results in a four hybrid sp3 atomic orbital

4 equivalent

H

bonds

Cl Cl

P Cl

s Cl Cl

11. 3 – Hybridization

Example #4: PCl5 (Trigonal bipyramidal)

5 hy

P s

p

d

Hybridization sp3d (5)

d

Results in a one sp3d hybridization

One of f sp3d hyb orbitals

Example #5: SF6 (Octahedral)

Results in a sp2d2 hybrid orbital

Hybridization in Atoms with Lone Pairs N But actual bond angle is 107°

Example: NH3

2s

N

2p

11. 3 – Hybridization sp3 hybridization!

s

p

Hybridization

N sp3

Hybrization in Atoms with Lone Pairs: Example #5: NH3 p orbitals are perpendicular N

N

N

2s

2s

s

2p

p

Hybridization

2p

N

Suggests 3 bonds

sp3

1lone pair

H H

Multiple Bonds:

C CC

bond angles 390° bonding pairs H H H

C C

C

H

c

1s

H 1s

c

1s

H H

C C

18 1lone pair C

sp3

C

C CC

sp3 C

H

1s

C

sp

p

sp 1s

sp3

3

bonds

p-p side by sid 1s

sp3

p orbitals are perpendicular C

H H

C CC

H H

C C

H

C

C

C

C

90° bond angles

c

H

C

C

C sp

18

p-p side by side

C

H