3. Simple Bonding Theory
3. 1 Lewis “Electron-Dot”
Diagrams & Resonance
3.1 Lewis-Representation of “simple” molecules such as
F2, N2, H2O, H-C≡C-H, or CO2 (but how about O2??)
3.1.1 Resonance Structures
3.1.2 Expanded Shells
1
3. 1.2 Expanded Shells
3.1.2 Expanded Shells
How about IF7 (14 electrons), [TaF8]3− (16 electrons)
…or even [XeF8]2− (18 electrons)
Is that it? And if so, why?
2 for s
6 for p
10 for d
+ steric crowding
[Mo(CN)8]4−
square antiprismatic
(COON 8)
[Re(H)9]2−
tricapped trigonal prismatic (COON 9)
very high
coordination
numbers COON
3. 1.2 Expanded Shells
3.1.3 Formal Charge
2
3. 1.2 Expanded Shells
& Formal Charges
3. 1.4 Multiple Bonds
in Be and B Compounds
+1
Multiple bonds for F and Cl
as well as B and Be
unreasonable;
unlikely formal charges
+1
-2
+1 for F, but resonance-stabilized
trigonal planar
prefers formation of adducts with
Lewis-bases such as H3N: Et2O: or THF
approx. tetrahedral
3
3. 2 VSEPR Theory
Valence-Shell-Electron-Pair-Repulsion Theory
described by Sidgwick & Powell (1940)
further developed by Gillespie & Nyholm (1957)
80 Years and still
VSEPRing
Ronald J. Gillespie (Aug. 21, 1924)
McMaster University
3. 2 VSEPR Theory
Required Skills:
Identify the central atom in a molecule containing more than two atoms as
a start.
Identify the number of valence electrons of any element. This concept is
important, because you need to know the number of valence electrons in
order to write a Lewis dot structure for the molecule in question.
Count the number of VSEPR pairs or steric number (SN) for the central atom
in a molecule. You need this number in order to describe or predict the shape
of the molecule in question.
Determine the number of lone electron pairs that are not shared with other
atoms. Often, a Lewis dot structure is useful to help you count this number.
Predict the shape of molecules or inos as the key concept of VSEPR theory.
From the shape and by applying the idea that lone electron pairs takes up
more space, you can predict the bond angles withing 5% of the observed
values.
Predict the values of bond angles and describe the hybrid orbitals used by
the central atoms in the molecules or ions.
4
3. 2 VSEPR Theory
Remember: Molecules adopt geometries in which their
valence electron pairs position themselves as far from
each other as possible!
AXmEn
Steric Number SN = n + m
A = central atom
X = any ligand
E = lone pair(s)
S of ligands & electron pairs!
EA
S
SN = 2
linear
Y!
SN = 3
trigonal planar
Multiple Bonding Does Not Effect the Geometry! Well, we’ll see…
3. 2 VSEPR Theory
5
3. 2 VSEPR Theory
Conversion of a
Cube into a
Square Antiprism
3. 2 VSEPR Theory
3.2.1 Lone Pair Repulsion
Differences between
lone pairs & bonding pairs
two
90o
interactions
three
90o
interactions
Assign the most space
to lone electron pairs!
6
3. 2 VSEPR Theory
3.2.1 Lone Pair Repulsion
OK, this one is tricky…
ax 169.8 pm
87.5o
eq
158.9 pm
…just follow the rules:
•90o
(or less, especially for lp-lp) interactions is the worst
• lp-bp interactions are more important to consider than bp-bp
3. 2 VSEPR Theory
3.2.1 Lone Pair Repulsion
…in tbp geometries,
lone pairs always
equatorial
and in octahedral
lp are axial
7
3. 2 VSEPR Theory
3.2.2 Multiple Bonds
Double and triple bonds
have slightly greater
repulsive effects than
single-bonds:
Tend to occupy the same
positions as lone pairs
(or compete):
3. 2 VSEPR Theory
3.2.3 Electronegativity: Scales, Geometry, and Size Effects
Electronegativity is a measure of an atom’s ability
to attract electrons from a neighboring atom
to which it is bonded!
H-H: EBE = 432 kJ/mol
Cl-Cl: EBE = 240
H-Cl: EBE(est) = 336
H-Cl: EBE(exp) = 436
8
3. 2 VSEPR Theory
3.2.3 Electronegativity: Scales, Geometry, and Size Effects
He & Ne do not form molecules (no known bonds): calculated from
ionization energies
Even more electronegative than F! Higher ionization, smaller radii
3. 2 VSEPR Theory
3.2.3 Bond Angles : Electronegativity or Does Size Matter ?
either way, works well!
doesn’t work at all
9
3. 2 VSEPR Theory
3.2.4 Bond Angles & Ligand-Close Packing Model
Similarly inconclusive!
3. 2 VSEPR Theory
Summary
Summary to predict the shape of a molecule:
•Write down the Lewis dot structure for the molecule
•Count the number of bond pairs and lone pairs around
the central atom
•Decide on the electron pair orientation for the total
number of electron pairs (4=tetrahedral,
6=octahedral…)
•Consider the placement of lone pairs and any distortions
from "regular" shapes
•Name the shape based on the location of atoms (nuclei)
10
3. 2 VSEPR Theory
Summary
3. 2 VSEPR Theory
Summary
11
3. 2 VSEPR Theory
Summary
3. 3 Polar Molecules
3. 3 Bond Dipoles & Molecular Dipoles
The polarity of the molecule is the sum of all of the bond
polarities in the molecule. Since the dipole moment (μ, measured
in Debyes (D)) is a vector (a quantity with both magnitude and
direction), the molecular dipole moment is the vector sum of the
individual dipole moments. If we compare the molecular dipole
moments of formaldehyde and carbon dioxide, both containing
a polar carbonyl (C=O) group, we find that formaldehyde is
highly polar while carbon dioxide is non-polar . Since CO2 is a
linear molecule, the dipoles cancel each other
12
3. 3 Polar Molecules
3. 3 Bond Dipoles & Molecular Dipoles
3. 4 Hydrogen Bonding
3.4 Hydrogen Bonding & Boiling Points
two lone pairs &
very polar O-H bonds
result in very large
net dipole
one lone pair &
polar N-H bonds
result in large
net dipole
sometimes opposing
polarities & difficult
to predict
13
3. 4 Hydrogen Bonding
3.4 The Effect of Hydrogen Bonding on Boiling Points
H2O: four H-bonds
HF: two H-bonds
”
mal
“nor ing MW
s
a
es
e
incr on forc
d
n
o
L
Now, think about all the awesome properties of “ice”…
3. 4 Hydrogen Bonding
3.4 The Effect of Hydrogen Bonding on “Life”
NH3
…but -N—H—O- is just fine…
…is tricky
-> secondary & tertiary protein structures
14