Formic acid: HCOOH
H
H C
Acetone
H
C
O
H
C H
H
Benzene C6H6
Kekulé structures
Resonance structures; each point corresponds to a CH
Each C is sp2 hybridized, one of the sp2 forming a s-bond
with H 1s orbital and the other two forming s-bonds with
adjacent C sp2 orbitals.
The un-hybridized p orbital on each C is available for pbonding with p orbitals on either of the adjacent C atoms
Actual structure of benzene is a resonance hybrid of the two
alternating bond patterns; the 6 C atoms are identical, and
the electrons in the p-bonds spread around the entire ring
This lowers the energy of the molecule - resonance adds
stability to a molecule
Characteristics of p bonds bonds
Energy of C=C is < 2 x energy of C-C bond
Energy of CC is < 3 x energy of C-C bond
C, N, O form double bonds with one another and with
elements from later periods
Double bonds are rarely found between elements in period
3 are below - atoms are too large for effective side-by-side
overlap.
Molecules with alternate double-single bonds - conjugated
molecules
Isomers: Molecules with the same molecular formula but
different structures
H
H
C
Cl
H
C
Cl
C
Cl
cis-1,2-dichloroethylene
Cl
C
H
trans-1,2-dichloroethylene
Rotation can occur about a single sigma bond
Rotation is restricted about a double bond; isomers are a
consequence
Change of shape triggers a signal along the optic nerve
Molecular Orbital Theory
VB theory: localized bond
VB theory provides the basis of calculating electron
distributions in molecules but cannot explain the
properties of some molecules.
O O
O 2:
VB theory
O: Is2 2s2 2p4
sp2 hybridized O, one sp2 from each forms s-bond and the
other two are occupied with the lone pairs.
The un-hybridized p on each forms the p-bond
Indicates that in O2 molecule, all electrons are paired.
However O2 was observed to be paramagnetic
VB theory assumes that the electrons are localized between
the two bonding atoms
Molecular orbital theory: electrons are spread throughout the
entire molecule; electrons are delocalized over the whole
molecule.
Pure atomic orbitals combine to produce molecular orbitals
that are spread out, delocalized, over an entire molecule
Molecular orbitals are built by adding together superimposing - atomic orbitals belonging to the valence
shell of the atoms in the molecules.
H2: wavefunction representing the molecular orbitals (MOs)
for H2 can be represented by combining the two atomic
orbitals (AOs) for the separated H atoms.
Wavefunction of the H2 MO
y+ = yA1s + yB1s
yA1s or yB1s 1s orbital centered on one of the H atom(A or B)
The molecular orbital, y, is a linear combination of atomic
orbitals
Any molecular orbital formed from a superposition of atomic
orbitals is called a LCAO-MO.
y+ is a bonding orbital; energy of y+ is lower than that of
either AO
In H2, the contribution from each AO to the MO is equal
The two AOS are waves centered on different nucleii.
Bonding orbital: AO wavefunctions interfere constructively MO wavefunction in blue.
N AOs overlapping will form N MOs
Two H AOs overlapping form two Mos; one of which is the
bonding orbital, y+.
The wavefunctions of the two H AOs can also interfere
destructively - anti-bonding MO of higher energy than each
of the AOs
y- = yA1s - yB1s
Node between two nuclei
Probability of finding electrons
between nuclei reduced; nuclei
repel each other
http://www.shef.ac.uk/chemistry/orbitron/index.html
Molecular Orbital Energy Level Diagram
Energy of bonding MO < AO
Energy of anti-bonding MO > AO
Diatomic Molecules
Build all possible MOs from available valence AOs
Then accommodate valence electrons in molecular orbitals
using the aufbau principles
1) Electrons occupy the lowest energy MOs first, then orbitals
of increasing energy
2) Pauli exclusion principle: each orbital can occupy up to
two electrons; if two electrons in an orbital must be paired
3) Hund’s rule: if more than one orbital of the same energy is
available electrons enter them singly with parallel spinds.
Lowest unoccupied MO
(LUMO)
Highest occupied MO
(HOMO)
H2 molecular orbital energy-level diagram or correlation
diagram
Bonding MO - s1s
anti-bonding MO s *1s
H2 Ground state electron configuration (s1s)2
Bond Order =
0.5(number of electrons in bonding MOs
- number of electrons in anti-bonding MOs)
H2+ bond order = 0.5
(s1s)1
H2 bond order = 1
(s1s)2
He2 bond order = 0
(s1s)2 (s*1s)2
Period 2 elements
In period 2 elements each atom has one 2s and three 2p
valence AOs; expect to form eight MOs
The two 2s orbitals (one from each atom) overlap to form a
s2s bonding MO and a s*2s antibonding MO
The six 2p orbitals (three from each atom) overlap to form six
MOs
The two 2p-orbitals directed toward each other form a
bonding s-orbital (s2p) and an anti-bonding s*-orbital (s*2p)
Two 2p orbitals that are perpendicular to the internuclear axis
overlap side by side to form two bonding p and two antibonding p* orbitals.
Antibonding
Bonding
s and s* orbitals
formed from p AOs
p and p* orbitals
formed from p AOs
MO diagram for
homonuclear diatomic
molecules Li2 through N2
MO diagram for
homonuclear diatomic
molecules O2 and F2