A brief review (chapter 1 in text)
Structure of atoms:
nucleus (protons, neutrons)
positive charge
“cloud” of electrons
negative charge
Opposites attract:
the closer electrons are to
nucleus, the lower their
energy (E).
Electrons reside in orbitals: region of space corresponding to a
high probability of finding a given electron.
1
Orbitals and shells
Each atom has several orbitals, Each with a distinct shape and size.
s-orbital
p-orbital
Orbitals are arranged in “shells” around the nucleus.
Each shell is a different distance from the nucleus, and so the electrons in the
shells have different energies (closer=lower E)
1st shell: 1 orbital (2 electrons)
2nd shell: 4 orbitals (8 electrons)
3rd shell: 4 orbitals (8 electrons)
Orbitals
1s
2s 2px2py2pz
3s 3px3py3pz
2
Filling Atomic Orbitals
1. Aufbau principle: orbitals are filled in order of increasing energy
2. Pauli Exclusion Principle: No more than two electrons may be present in an orbital,
one with spin number +1/2 and the other with spin number -1/2.
3. Hund’s rule: When orbitals of equivalent energy are available but there are not
enough electrons to fill all of them completely, then one electron is added to each orbital
before a second is added to any one of them.
3
Atomic Orbitals of Carbon
6
C
Carbon has 6 electrons
12.011
2s
Energy
2p
1s
4
Molecular Orbitals
Molecular orbitals (MOs) are formed by the combination of 2 atomic orbitals (AOs). These
can be combined by addition of the AOs to form a bonding orbital or by subtraction of one
orbital from the other to form an antibonding orbital.
Bonding orbital: lower in energy
relative to atomic orbitals that
combined to form it.
Antibonding orbital: higher in energy
relative to atomic orbitals that
combined to form it.
e.g. 2 H
*
H2
5
Carbon Molecules
C
2py
2x
H
2px
1s
2p
C
x
90°
H
2s
1s
H
“carbene”: highly
unstable (reactive)
Only 1/2 filled AOs form combine to form MOs: predicts C would
form only 2 bonds. This is not what we see!
6
Hybridization
Hybridize 2s and 2p orbitals to form 4 sp3 orbitals.
2p
3
2sp
2s
1s
1s
3x 2p
4x
1x
2s
2sp3
7
What we see:
The 4 sp3 orbitals of carbon can interact with the 1s orbitals of hydrogen or the
sp3 orbitals of another carbon to form new σ-bonds (sigma bonds)
H
C
4x
H
H
C
H
H
methane
σ bonds
2x
C
6x
H
H
H
H
C
C
H
H
ethane
H
8
3D Structure of Molecules
H
H
C
C
H
tetrahedron
H
109.5°
Methane
H
H
H
H
In order to minimize the repulsion between electrons, the 4 sp3 orbitals are
oriented as far away from each other as possible. This is accomplished by the
orbitals taking on a tetrahedral geometry around the central carbon atom. As a
result, the hydrogens in methane and other sp3 hybridized carbons reside on the
vertices of a tetrahedron, with the carbon in the center.
9
Structural Notation
in the plane of
board/paper
H
in the plane of
board/paper
in the plane of
board/paper
bond “going back”
C
H
H
behind the plane
of board/paper
H
bond “coming out”
in front of the
plane of
board/paper
10
Nitrogen and oxygen also hybridize
lone pair
electrons
bonding
electrons
lone pair
sp3 N
ammonia
N
H
H
H
107.3°
lone pair
electrons
sp3 O
bonding
electrons
lone pairs
water
O
H
H
104.5°
11
sp2 hybridization
Hybridize 2s and 2p to form 3 sp2 orbitals and 1 p orbital.
2p
2p
2
2sp
2s
1s
atomic orbitals
1s
hybrid orbitals
Now have 3 orbitals to form sigma bonds with: 3 identical sp2 orbitals.
The electron in 2p-orbital does not form σ-bonds.
12
sp2-hybridization and geometry
2p
2
2sp
1s
C
trigonal planar
(equalateral triangle)
120°
σ bonds
H
H
C
H
C
120°
H
13
What about the left over π-orbitals?
π* -orbital
anti-bonding (high energy)
2p
π-orbital
bonding (low energy)
H
H
C C
H
H
double bond
14
sp2-hybridized nitrogen and oxygen
2p
2
2sp
H
••
N
N
C
H
1s
H
2p
••
••
O
2sp2
C
O
H
1s
H
formaldehyde
15
sp-hybridization
2sp3
x
2p
1s
2s
2p
2sp
1s
Berylium
AOs
1s
16
Geometry of sp-hybridized structures
180°
F
Be
F
sp-hybridized systems are linear
17
Carbon and sp-hybridization
hybridize 2s and 1 x 2p to form 2 x sp orbitals and 2 x p orbitals.
2p
2p
2s
1s
2sp
1s
Now have 2 orbitals to form sigma bonds with: identical sp orbitals point
in opposite directions to each other.
18
Sigma framework
π-bonds
p-orbital
(y-direction)
H
•
•
•
C
•
C
H
C
C
H
H
C
C
C
H
C
H
H
H
p-orbital
(z-direction)
π-bond
(z-direction)
π-bond
(y-direction)
σ-bonds
(x-direction)
H C C H
triple bond
4x p AOs combine to form 2x π-bonds
19
sp-hybridized nitrogen (but not oxygen)
2p
N
2sp2
C
N
cyanide
1s
O
2p
2sp2
This leaves no hybridized
orbitals to form σ-bonds!
1s
20
Atoms are hybridized not molecules!
sp2
sp2 O
H
H
C
C
C
C
H
H
H
sp2
C
C
sp3
sp2
sp
H
sp
21
Overview of
possibilities
22