Writing Lewis Structures
-keeps count of valence electrons
(electrons in the outer shell - participate in bonding)
Rules:
1. Every atom would like to have or share 8 valence e-’s
Exceptions: H and He would like to have only 2 e-’s
3rd and 4th row elements can break this rule. (d-orbitals)
2. Each atom must have (not share) the correct # of
valence e-’s or else a charge must be assigned.
re
b
Ab
r
o
f
v.
.
m
e
h
C
.
g
r
O
Formal Charge
-Often convenient to associate formal charge with an atom
Formal Charge =
# of
valence
e-’s
-
# of
unshared
e-’s
Or bonding that differs from ‘normal’
Neutral
Charged
-
1/2 the # of
shared
e-’s
Formal Charge - How realistic?
N-C
Are the electrons equally shared?
Electronegativity: ability of an atom to attract electrons towards itself
in a covalent bond
http://en.wikipedia.org/wiki/Electronegativity
Other trends we should review?
molecular
atomic
Pauling definition: “the power of an atom in a molecule to
attract electrons to itself.” (calculate by bond dissociation
energies)
Mulliken definition: average of ionization potential of an
atom and its electron affinity
So how equally are electrons shared?
Covalent: ΔEN < 0.5
Polar Covalent ΔEN between 0.5 to 1.9
Ionic ΔEN > 1.9
In a single N-C bond
3.0 - 2.5 = 0.5
Polar
Electrostatic Surface Potential: In complex molecules we calculate
the electrostatic potentials or electron density (Spartan, ChemDraw,
freeware, Macromodel, etc)
Polar Bonds
Dipole
Electrostatic surface potentials
van der Waals surfaces
Energy (kJ/mol)
Computed with 6-311++G** basis set
Electrostatic surface potentials
Who will participate in hydrogen bonding?
Acceptor - A:
Donor - D-H
Polarizability α (cm3/10-24)
temporary dipole within atom or molecule (microscopic timescale)
ability of an electron cloud to be perturbed
(correlates with atomic volume)
how well bond reacts to exterior charge
Inductive Effects:
Withdrawing electrons through σ-bonds
+ 0.09
+ 0.36
δ−
- 0.02
- 0.17
Magnitude of the delta charges
diminishes but alternates
Field Effects - polarization in a molecule that
results from charges that interact through space.
Bond dipoles, Molecular dipoles, quadrupoles:
Need information on the Shape of the molecule:
Valence Shell Electron Pair Repulsion rule (VSEPR)
Count the number of ‘groups’ around an atom
Counts as one ‘group’
single bond
double bond
triple bond
lone pair
# of groups
2
3
4
geometry
linear
trigonal planar
tetrahedral
bond angle
180°
120 °
109.5°
Reality Check
•How do these compare?
Molecular dipole moment µ (Debye D)
•vector sum of the individual bond dipoles
BF3
CCl4
µ=0
µ=qxr
q = charge and r is distance
HCN
µ= 2.33 D
µ 1 D = 10-18 electrostatic units, the charge of one e-
H 2O
Resonance - more than 1 good Lewis structure
Some compounds can be assigned more than one good Lewis structure.
The compound behaves experimentally like hybrids or molecular averages of
the individual Lewis Structures. These molecules are called resonance
hybrids and only differ by movement of electrons.
Note: Resonance always stabilizes a structure.
Resonance
When is a particular Lewis Structure a significant
contributor to a resonance hybrid??
1. Same or nearly the same number of bonds.
(There might be a structure with one less bond if it has
more reasonable formal charges)
2. Formal charges on reasonable atoms. e.g.
Reasonable:
Less reasonable:
NOT reasonable:
Much less
reasonable:
Things to Check:
1) Charge conservation on either sides of an arrow
2) Octet Rule (do any of the atoms exceed an octet?)
3) Electrons moving in the right direction (toward charges
or δ+ charges)
4) Total number of valence electron should not change
Evaluating Resonance Structures:
1) Minimal Charge separation - The best situation is with
no charge separation (ie no overall charge)
2) Full Octets (or Duet for Hydrogen)
3) Charges on atoms of appropriate electronegativity
- Negative charges on more electronegative atoms
( Cl > O > N > C )
- Positive charges on less electronegative atoms
( C > N > O > >Cl )
4) More bonds the better.
* Note: sometimes these 'rules' will conflict.
Let’s evaluate some structures….
Who is more stable?
Diazomethane: CH2N2
1.
H
N
N
N
C
H
2.
H
N
C
H
N
H
N
C
H
Let’s evaluate some structures….
Rank from most stable to least stable
A.
B.
Valence Bond Theory (VBT)
Good Overlap
Atomic orbitals
overlap to form a
chemical bond.
+
s
s
+
The electrons are
localized in the bond
region.
Orbitals must have the
correct symmetry (+/+)
or (-/-) to overlap
s
px
py
pz
px
+
sigma bonds (σ)
end to end overlap
px
py
pz
pi bonds (p)
sideways overlap
Most orbital pictures are taken
from MIT CHEM 5.12 open
courseware
Valence Bond Theory (VBT)
Out of Phase (+/-): No Overlap
s
px
Orthogonal: No Overlap
s
py
px
px
s
pz
px
py
py
py
pz
pz
px
px
pz
pz
Atomic Orbitals and Bond Angles
VSEPR Theory - # of e- domains surrounding an atom dictates its geometry
Single,
double, triple
bond or lone
pair
4 e- domains
3 e- domians
2 e- domains
180o
109.5o
120o
X
linear
trigonal
planar
tetrahedral
But the e- density in atomic orbitals are orthogonally distributed
All angles in all molecules would be 90o!
y
y
x
x
z
y
z
x
z
How is this geometric mismatch explained?
Hybridization sp3
C
C:
s
p
4 sp3
4 edomains
25% s - 75% p 109.5o
+
+
+
Electron density distributed to the
corners of a tetrahedron
4 atomic orbitals
2s + 2px + 2py + 2pz
yield
4 hybrid atomic orbitals
4 sp3 orbitals
Hybridization sp3
Methane
Ethane
H
H
H H
H C H
C H
H
H
H C C H
H
s-sp3 σ bond
HH
H
C
HH
H H
sp3-sp3 s bond
single bonds are σ bonds
C
H
Hybridization sp2
+
C:
C
3 edomains
s
sp2
p
p
33% s - 67% p
120o
+
+
4 atomic orbitals
{2s + 2px + 2pz }, 2py
sp2 electron density distributed
to the corners of a triangle
(trigonal planar)
yield
3 sp2 hybrid orbitals
+ 1 py
Hybridization sp2
Ethylene
H
H
C C
H
H
py-py π bond
3 sp2 + 1 py
sp2-sp2 σ bond
double bond = sigma bond (σ) + pi (π) bond
Let’s evaluate some structures….
Rank stability
A.
B.
C(t-Bu)3
C(CF3)3
Hybridization sp
C
2 edomains
+
C:
s
sp
p
p
50% s - 50% p
180o
X
+
sp electron density
distributed linearly
4 atomic orbitals
{2s + 2px}, 2py, 2pz
yield
2 sp hybrid orbitals
+ 1 py + 1 pz
Hybridization sp
Acetylene
H C C H
py-py p bond
2 sp + 1 py + 1 pz
pz-pz p bond
sp-sp s bond
triple bond = sigma bond (σ) + 2 pi (π) bonds
Molecular Orbital Theory
Two ways to describe bonding:
1) Valence Bond Theory (VB Theory, already discussed)
2) Molecular Orbital Theory (MO Theory)
A rigorous, mathematical description of bonding
(we’ll do a simple, qualitative examination)
Two atomic orbitals on different atoms combine to form two molecular orbitals:
1. A low-energy bonding orbital
2. A high-energy antibonding orbital
VB Theory - overlap of atomic orbitals, e- density localized in bond
MO Theory - combination of atomic orbitals to form molecular orbitals, edensity spread throughout entire molecule