Important concepts:
Octet rule
Formal charge
Resonance
Shapes of molecules (VSEPR)
Isoelectronic principle
Bond polarity and bond strength
Lewis dot structures: based on stable s2p6 (octet) configuration
H
H
:N H
H
:
:
:N : H
H
=
Formal charge
To get formal charge, divide bonds equally between atoms.
E.g., NH3
N has 5 e} both atoms are neutral
H has 1 e-
H + - F
H N B F
H
F
H
:O H
H
+
-
-
O
O
S O-
2+
O-
Formal charge is NOT the same as the oxidation state or
the actual charge on the atom.
Formal charges must add up to overall charge on molecule.
Formal charge
Low formal charges are most stable.
Can be used to tell if a proposed structure is correct or likely.
F
E.g., BF3
B
F
F
F
F
Non-octet
+
F
N
O
B
+
-
Octet form is unlikely,
because + charge is on
F, which is more
electronegative than B.
-
F
Octet
vs.
O
N
F
More reasonable
Resonance Structures
Ozone
-
O
+
+
O
O
O
O
O
-
Terminal O atoms are spectroscopically equivalent
Real (instantaneous) structure is average of two forms:
+
-1/2
O
O
O
-1/2
Resonance Structures - another example
O
Nitrate ion
O
N
-
O
+
O
-
+ 2 other forms
O
all O’s equivalent
N
+
O
-2/3
bond order 1.33
Resonance Structures: Non-equivalent structures
OCN(cyanate ion)
-O
C
N
O
C
Which of these is not an
important resonance structure?
N-
+O
C
N2-
• High formal charge
• O is more electronegative than N
Real electronic distribution is a weighted average
of the first two non-equivalent resonance structures
Resonance Structures: No-Bond Resonance
X
Y
+X
Z
Y
Z-
ONF3 Electron diffraction shows unusually long N-F bond. Why?
-O
F
N F
+
+
O N
F
F
H H B
BH3 + CO
Lewis
acid
F
F- + 2 other equivalent forms
Lewis
base
C O
H+ H B C O
+
H (long B-H bond)
H
Resonance Structures: Hypervalent Compounds
e.g.,
I- + I2 ! I3-
hypervalent structure:
triiodide ion
.. ....- ..
:I.. - ..I - ..I:
central I has 10 e- (violates octet rule)
Vibrational spectra show I-I bond is weaker than in I2
I
I
I-
I
-
I
I =
-1/2I
.. .. octet
-1/2
..I I
bond order = 1/2
Resonance Structures: Hypervalent Compounds
XeF2: Same number of valence electrons as I3+
-1/2F
F-1/2
Xe
XPS shows partial (-) charge on F
IR shows long (weak) Xe-F bond
relative to XeF+
+
Xe -F single bond
(like I2)
Resonance Structures: Hypervalent Compounds
F
SF6:
F
F
S
F
F
F
12 e- around S
F
F
F
F
F
F-
S2+
F
F-
F
+ other
resonance
structures
F
S2+
F
F-1/3
bond order = 2/3
Resonance Structures: Hypervalent Compounds
Valence shell d-orbitals can be used to create “expanded”
octets for 3rd and higher periods:
F
F
F
S
F
O-
O-
O S O
O Cl O
O-
O-
F
F
+
3d orbitals are high in energy
They participate somewhat in bonding, but most molecular
properties (shape, bond orders…) are adequately explained
using only s and p hybrids.
The Isoelectronic Principle
Molecules and ions with the same number of valence
electrons have similar geometries and properties
e.g.,
O
H2C
C O
CO2
C
allene
CH2
-
-
+
N
N O
N2O
N
N
+
N
-
azide ion
All four have 3 heavy atoms, 16 valence electrons
The Isoelectronic Principle
Which one of these does not belong??
F
F
F
F
O
B
N
C
N
F
F
F
O
F
O
-
O
+
O
-
O
-
C
O
-
Which of the following is (are) isoelectronic with H2O??
CH4
F-
NH3
OH-
BH4-
The Isoelectronic Principle works for solids too
Al Si P
S
Cd Ga Ge As Se
In Sn Sb Te
CdSe, GaAs, Ge are all
isoelectronic semiconductors
InP, AlSb are similar too
AlPO4 - same structures as SiO2
similar physical properties
Electron pair repulsions can be used to rationalize molecular
shapes (Valence Shell Electron Pair Repulsion Theory)
Rules:
Lone pairs + atoms define “total coord. no” of a central atom.
Electron pairs + bonds orient in space to minimize repulsion
Repulsive interactions are lp-lp > lp-b > b-b
Total coord #
2
3
Total coord #
Shape
180o
linear
e.g., CO2
triangular
e.g., NO3-
120o
Shape
109.5o
4
tetrahedral
5
trigonal bipyramidal
(more common)
CH4
90o
equatorial
120o
PF5
axial
square pyramidal
(less common)
Sb(Ph)5
90o
6
octahedral
SiF62-
C
:
O
:
O
:
Examples:
:
2 bonds to central atom, no lone pairs
Total coord. # = 2 (linear)
2 bonding domains + 1 lone pair
:
SO2
S
O
+
Electronic shape is trigonal (CN = 3)
O
-
Molecular shape is bent
O-S-O angle is slightly less than 120o
(lone pair repulsion)
H
104.5o
H2O
:
H
Electronic shape: tetrahedral
:
Molecular shape: bent
-
Examples:
-
O
O
S O-
2+
Tetrahedral, no lone pairs on S
F-
"
F-
:
:
F
:
Bond order = 1/2
:
F -1/2 --- Xe+ --- F-1/2
Total C.N. = 5
:
Xe+ - F
:
:
F - Xe+
:
:
XeF2
:
O-
:
Xe
:
F
Axial bond order = s1/5pz1/2 = 0.7 (formal charge = -0.3)
Equatorial bond order = s1/5px,y2/3 = 0.867 (f.c. = -0.133)
(equatorial bonds are less ionic)
More electronegative ligands (F)
go to axial sites in TBP, so
molecule is linear.
F
:
Br+
:
F
BrF4-
F-
+ other resonance forms
F-
:
F
Br+
:
F
F
F
Formal charge = -1/2
:
:
Bond order = 1/2
Is geometry
or
:
:
?
Molecule is square planar
Minimizes lp-lp repulsion
1.68 Å
:
BrF5
Show that no-bond resonance predicts
shorter axial than equatorial bonds
1.78 Å
Pauling introduced the concept of electronegativity (#) to
explain the extra bond energy of polar molecules
H-H
H-H
O=O
H
H
A-A + B-B
Homonuclear diatomics
or single bond in more
complex molecule
O
O
H
H
!+ !2 A-B
Very exothermic, but
number of bonds is the
same (4) on each side.
Extra bond energy from
electrostatic attraction
Bond energies for single bonds are E(AA), E(BB), and E(AB)
E(AB) = 1/2[E(AA) + E(BB)] + 23 (# A - # B)2 (kcal/mol)
Pauling scale for # is 0.7 (Fr) to 4.0 (F)
If $ # > 2.0, the bond is ionic
If 0.5 < $ # < 2.0, the bond is polar covalent
And if $ # < 0.5, the bond is non-polar.
Linus Pauling
The polarity of bonds explains something about reactivity:
!- !+
!+ !e.g. Si - H bond is more hydride-like than C - H
1.8 2.1
2.5 2.1
So silanes react with strong acids to make H2, but phosphines and C-H
compounds do not.
Electrophilic substiution reactions occur easily on Si - H, P - H compounds
2.1 2.1
For many bonds, Pauling’s formula is obeyed:
D(n) = D(1) - 0.6 log(n)
D(1) = single bond length (in Å), D(n) = length of bond order n
e.g., C-C bond in alkanes is 1.54 Å
C=C (ethane)
1.36
1.33
C%C (acetylene)
1.25 Å predicted
1.20
observed
Some bond energies and bond lengths are anomalous
e.g., F-F bond in F2
F covalent radius is 0.64 Å, but F-F bond length is 1.43 Å (extra 0.15Å)
calc’d bond order 0.6
X
F
Cl
Br
I
E(XX) (kcal/mol)
38
Low bond dissociation energy makes F2
58
a very powerful oxidizing agent
46
(oxidizes Xe, Kr, O2 )
36
.. ..
:F - F:
.. ..
1s2 core
Lone pair repulsion stretches F-F bond
Cl2 is “normal” because core is larger
covalent radius = 0.99 Å
bond distance = 1.98 Å