CH2. Molecules and covalent
bonding
Lewis Structures
VSEPR
MO Theory
1
Lewis structure H3PO4
•
Skeleton is:
•
Count total valence electrons:
1P= 5
3H= 3
4 O = 24
Total = 32 e- or 16 valence e- pairs.
•
7 e- pairs needed to form s skeleton.
2
Lewis structure H3PO4
• Add remaining e- pairs:
• Left has a formal charge of +1 on P and -1 on one O,
right has 5 e- pairs around P (hypervalence)
• Analysis of phosphoric acid shows purely Td phosphate
groups, which requires something beyond either simple
Lewis model.
3
Resonance in NO3-
experimental
data - nitrate is
planar with 3
equivalent N-O
bonds
4
VSEPR model
• Count e- pairs about the central atom (draw
Lewis structure if needed). Include non-bonding
pairs, but not multiple bonds.
• Geometry maximizes separation:
# e pairs geometry
example
2
3
4
5
6
7
8
linear
equilateral triangular
tetrahedral (Td)
trigonal bipyramidal (TBP)
octahedral (Oh)
pentagonal bipyramidal
square antiprismatic
HF2BF3
CF4
PF5
SF6
IF7
TaF835
Drawing Oh and Td molecules
It's often useful to draw octahedra and
tetrahedra with a cubic framework
6
Deviations from ideal geometries:
unshared pairs and multiple bonds require larger bite
ex: CH4, NH3, H2O
<H-C-H = 109.5°,
<H-N-H = 107.3,
<H-O-H = 104.5
ex: ICl4-
6 e pairs around I, 2 lone pairs and 4 e pair bonds to Cl
Oh coordination, and geometry is square planar (lone
pairs are trans, not cis)
7
POCl3
note that in :PCl3
the <ClPCl = 100.3,
the lone pair is
more repulsive
towards other
ligands than the
multiple bond !
based on Td geometry
< ClPCl = 103.3° due to
repulsion by multiple bond
Ligands move away from multiple bond
8
XeF5+
5 Xe-F bonds and 1 lone
pair on Xe
geometry based on Oh
coordination
lone pair repulsion gives
< FeqXeFeq = 87°
< FaxXeFeq = 78°
9
Fajan’s rule
bond polarization is towards
ligands with higher c,
decreasing repulsive effect.
Lone pairs are the most
repulsive.
ex: NH3 vs NF3
< HNH = 107.3°
< FNF = 102.1°
10
Inert pair effect
•
In Sn Sb Te
Tl Pb Bi
•
VSEPR geometries require
hybridization (valence bond term) or
linear combinations (MO term) of
central atom orbitals. For example,
Td angles require sp3 hybrid orbitals.
More on this in MO theory section.
Period 5 and 6 p-block central atoms
often show little hybridization (ex:
they form bond with orbitals oriented
at 90° as in purely p orbitals). This
can be ascribed to the weaker
bonding of larger atoms to ligands.
11
Inert pair effect - evidence
•
Bond angles near 90°:
NH3 107.2
AsH3 91.8
SbH3 91.3
•
PbO unit cell
•
H2O
H2Se
H2Te
104.5
91
89.5
Increased stability of lower oxidation
states
ex: Si, and Ge are generally 4+, but Sn
and Pb are common as 2+ ions (as in
stannous fluoride SnF2)
ex: In and Tl both form monochlorides,
B, Al, Ga form trichlorides.
Vacant coordination sites where the
lone pair resides
ex: PbO
12
Fluxionality
•
•
•
•
•
PF5 if TBP has 2 types of F ligands (equatorial
and axial).
19F NMR spectra at RT show only a single peak
(slightly broadened).
PF5 is fluxional at RT, i.e. the F ligands exchange
rapidly, only a single "average" F ligand is seen
by NMR.
Only occurs if ligand exchange is faster than the
analytical method. IR and Raman have shorter
interaction times and show 2 types of P-F
bonding at RT.
Even low temp NMR studies cannot resolve two F
environments
13
Berry pseudo-rotation
Sequences of the MD-Simulation of PF5 at 750K
(Daul, C., et al, Non-empirical dynamical DFT calculation
of the Berry pseudorotation of PF5, Chem. Phys. Lett.
1996, 262, 74)
14
Molecular Orbitals




Use linear combinations of atomic orbitals to
derive symmetry-adapted linear combinations
(SALCs).
Use symmetry to determine orbital interactions.
Provide a qualitative MO diagram for simple
molecules.
Read and analyze an MO diagram by sketching
MO’s / LCAO’s, describing the geometric affect
on relative MO energies.
15
H2
16
Some rules



The number of AO’s and MO’s must be equal.
This follows from the mathematics of independent
linear combinations.
More on symmetry labels later, but they come
from the irreducible representations for the point
group. s MO’s are symmetric about bond axis, p
MO’s are not. Subscipt g is gerade (has center of
symmetry), u is ungerade. Antibonding orbitals
are often given a * superscript.
The bond order = ½ (bonding e- - antibonding e-).
The bond energy actually depends on the
energies of the filled MO’s relative to filled AO’s.
17
O2
• MO theory predicts 2 unpaired e-,
this is confirmed by experiment.
• Bond order = ½ (8-4) = 2, as in
Lewis structure.
• MO indicates distribution and
relative energies of the MO's, Lewis
structure says only bonding or nonbonding.
18
I and Ea for atoms and diatomics
species I (kJ/mol)
N
O
O2
NO
F
F2
C
C2
1402
1314
1165
893
1681
1515
1086
Ea
142
43
123
300
19
Li2 – F2 MO’s
20
Some diatomic bond data
bond order
r0 in pm
D0 in kJ/mol
O2
2
121
494
O2 -
1½
126
O22-
1
149
F2
O2+
NO
NO+
N2
1
2½
2½
3
3
142
112
115
106
110
155
942
21
Spectroscopic data for MO’s
22
HF
23
Ketalaar triangle
HF
24
Hybridization
• Linear combinations of AO’s from same atom
makes hybrid orbitals.
• Hybridization can be included in the MO
diagram.
• In MO theory, any proportion of s and p can
be mixed (the coefficients of the AO’s are
variable). sp and sp3 hybrids are specific
examples.
25
H3+
26
BeH2
27
Correlation diagram for MH2
M
< HMH
Be
180°
B
131
C
136
N
103
O
105
28
Bonding MO’s in H2O
29
NH3
Use triangular H3 MO’s from
above as SALC's of the H
ligand orbitals. Must relabel
to conform with lower
symmetry pt group C3v. They
become a1 and e.
Combine with N valence
orbitals with same symmetry.
30
NH3 --calculated MO diagram
31
SF6
See textbook
Resource Section 5
for SALCs
32