Bonding
(Textbook Chapter 6 & 7)
Bonding (Textbook Chapter 6 & 7)
1. Electronegativity 2. Lewis Structures
3. VSEPR Theory 4. Orbital Overlap Model
5. Hybrid Orbital Theory
6. Molecular Orbital Theory
Bonding (1. Electronegativity)
When determining the structure of a chemical
compound, we pay close attention to
Attraction and Repulsion
Goal: arrange charged particles to minimize repulsion
– e- and nuclei attract (releasing E, stable)
– e- repel other e- (reduces stability)
– nuclei repel other nuclei (reduces stability)
Solution: concentration of e- between nuclei
Bonding (1. Electronegativity)
When electrons are in the region between two nuclei,
Attractive energies > Repulsive energies
(e-/n) > (e-/e- or n/n)
.
chemical bond
e- are not point charges but
are spread out over a large volume.
Olmsted, Chemistry, 3e, Canadian Edition
Bonding (1. Electronegativity)
A stable molecule is achieved when
Interaction Energy (J)
Attractive energies > Repulsive energies
(e-/n) > (e-/e- or n/n)
Inter-nuclear separation (pm)
Bonding (1. Electronegativity)
The interaction energy of two hydrogen atoms depends on
the distance between the nuclei.
Olmsted, Chemistry, 3e, Canadian Edition
Bonding (1. Electronegativity)
Some stable bonds share electrons evenly, but some
bonds have uneven electron sharing:
Electronegativity (Χ):
– an atom’s ability to attract bonding e(larger ΔΧ, more polar bond)
Ionization Energy (I1):
– how strongly an atom attracts its own e(how easy it is to remove that e- and ionize the atom)
Electron Affinity (EA):
– how strongly an atom attracts a free e-
Bonding (1. Electronegativity)
Electronegativity
Increase
Increase
Bonding (1. Electronegativity)
Electronegativity difference (ΔΧ) determines
the type of bond between two atoms:
B
A
covalent
0
0.5
polar covalent
ionic
2.0
Electronegativity difference (ΔΧ = ΧA – ΧB )
3.5
Bonding (1. Electronegativity)
Electronegativity
Examples:
Cl-Cl
H-Cl
(XCl = 3.0)
(XCl = 3.0)
ΔΧ = 0.0
covalent
(XCl = 3.0)
(XH = 2.1)
ΔΧ = 0.9
polar covalent
Bonding (1. Electronegativity)
Electronegativity
More Examples:
H-O-H
NaCl
(XH = 2.1)
(XO = 3.5)
ΔΧ = 1.4
polar covalent
(XNa = 0.9)
(XCl = 3.0)
ΔΧ = 2.1
ionic
Bonding (1. Electronegativity)
Electronegativity (visualizing e- sharing)
Cl
Cl
H
Cl
Na+ Cl-
Covalent
Polar Covalent
Ionic
e- equally shared
e- unequally shared
cation + anion
(e- transfer)
Bonding (1. Electronegativity)
Covalent Bonding
Interaction Energy (J)
Cl
e-/n attraction dominates
Inter-nuclear separation (pm)
Cl
Bonding (2. Lewis Structures)
Lewis Structures
– Schematic drawings of molecules
– Show distribution of bonding e-, valence e- and
nuclei (atoms) in molecules
• WHY bother? use to predict molecular shapes
– Molecular shape leads to further chemistry
understanding
Bonding (2. Lewis Structures)
How to build Lewis structures
1. Count valence e- for each atom (group #) and total
valence e- for molecule.
2. Put least electronegative atom in the middle.
3. Connect to outer atoms with bonds (2 e- per bond)
4. Assign remaining e- to outer atoms as lone e- pairs
to complete terminal atom octets. Add remaining
LPs to center atom.
5. Change LPs to bonds if necessary to ensure octets
where necessary
Bonding (2. Lewis Structures)
Lewis Structures: Example PCl3
1. Count valence e- for each atom and total valence
for molecule
P:
Cl:
1 x 5e- = 5e3 x 7e- = 21e26e-
Bonding (2. Lewis Structures)
Lewis Structures: Example PCl3
2. Put least electronegative atom in the middle
3. Connect to outer atoms with bonds (2 e- per bond)
Cl
P
Cl
Cl
(26 e-) – (3 bonds x 2 e-/bond) → 20 e- remaining
Bonding (2. Lewis Structures)
Lewis Structures: Example PCl3
4. Assign remaining e- to outer atoms as lone e- pairs
to complete terminal atom octets…
··
··
P
Cl
··
··
Cl
··
··
··
··
Cl
··
(20 e-) – (9 LP x 2 e-/LP) → 2 e- remaining
Bonding (2. Lewis Structures)
Lewis Structures: Example PCl3
4. Assign remaining e- to outer atoms as lone e- pairs
to complete terminal atom octets. Add remaining LPs
to centre atom.
··
P
··
Cl
··
··
··
Cl
··
··
··
··
Cl
··
(2 e-) – (1 LP x 2 e-/LP) → 0 e- remaining
Bonding (2. Lewis Structures)
Lewis Structures: Example CH2O
1-4:
CH2O (formaldehyde) has 12e-
H
6eC
2×H = 2×1 eO = 6 eC = 4 e-
H
··
··
O
··
…but C doesn’t yet have a full octet
Bonding (2. Lewis Structures)
Lewis Structures: Example CH2O
5. Change LPs to bonds if necessary to ensure octets
where required
H
C
H
··
··
O
··
…but C doesn’t yet have a full octet
Bonding (2. Lewis Structures)
Lewis Structures: Example CH2O
5. Change LPs to bonds if necessary to ensure octets
where required
H
C
H
··
··
O
… now it does.
Bonding (2. Lewis Structures)
Lewis Structures: Verifying a likely structure using
Formal Charge
Formal charge = [# of valence electrons] – [electrons
in lone pairs + 1/2 the number of bonding electrons]
For a neutral molecule, formal charge = 0
Formal charge = [# of valence electrons] –
[electrons in lone pairs + 1/2 the number of
bonding electrons]
Bonding (2. Lewis Structures)
Formal charge = [# of valence electrons] – [electrons in lone
pairs + 1/2 the number of bonding electrons]
FCP = 5 – [(3x1)+2] = 0
FCCl = 7 – [(1x1)+6] = 0
FCPCl3 = FCP + (3xFCCl) = 0
FCC = 4 – (4x1) = 0
FCH = 1 – (1x1) = 0
FCO = 6 – [(2x1)+4] = 0
FCCH2O = FCC + (2xFCH) + FCO = 0
Bonding (2. Lewis Structures)
Some molecules have more than one likely structure:
Example: CS2
both C and S have X = 2.5, so should structure be
S=C=S
OR
C=S=S
?
Bonding (2. Lewis Structures)
Resonance Structures
Example: CS2
Using the steps we’ve learned:
··
··
Valence ee- assigned
Formal Charge
··
··
C=S=S
C
4
S
6
S
6
6
-2
4
-2
6
0
Technically, this is Fctotal = 0,
but it is not the lowest integers to get to a sum of 0
Bonding (2. Lewis Structures)
Example: CS2
Using the steps we’ve learned:
S=C=S
··
Valence ee- assigned
Formal Charge
··
··
··
S
6
C
4
S
6
6
0
4
6
0
0
This is the more likely structure.
Bonding (2. Lewis Structures)
Example:
ClO3Cl: 1 x 7 e- = 7 eO:
3 x 6 e- = 18 e-:
1 x 1 e- = 1 e26 eFCCl = 7 – (5) = +2
FCO = 6 – (7) = -1
Bonding (2. Lewis Structures)
Bonding (2. Lewis Structures)
The position of the atoms is the same in the various
Resonance structures of a compound, but the position of
the electrons is different.
Minimize the formal charge on each atom
OR
Bonding (2. Lewis Structures)
Group 3A (13) compounds break the octet rules
Examples: AlCl3, BF3, trimethylaluminum…
··
··
··
··
F
··
B
F
··
··
Cl
··
··
··
F
··
Al
··
Cl
··
··
··
··
Cl
··
··
H3C
Al
CH3
CH3
Incomplete octets cause theses MX3 species to be Lewis acids
Bonding (3. VSEPR Theory)
We have a lot of information about a compound
without the structure…
Example:
Chemical Formula:
C21H30O2
Molecular Mass:
314.47g/mol
Compound Name: tetrahydrocannabinol
Bonding (3. VSEPR Theory)
We have a lot of information about a compound
without the structure…
…so why do we care about molecular shape?
Example: tetrahydrocannabinol
A.K.A. THC
Its structure has similarities to other known psychedelics,
so we can predict properties using molecular shape!
Bonding (3. VSEPR Theory)
Why we care about molecular shape:
• Chemical names and formulae don’t tell us how the
molecule really “looks”
– and therefore behaves
Consider enzymes and substrates:
+
=
Substrate pocket
→ enzyme shape
→ supramolecular shape
→ constituents’ molecular shapes
Bonding (3. VSEPR Theory)
Lewis structure of methane.
Modern experiments shows a tetrahedral structure
Bonding (3. VSEPR Theory)
To predict molecular shapes, we will use
VSEPR Theory
Valence Shell Electron Pair Repulsion
How?
– Draw Lewis Structure
– Count the number of electron groups around the
central atom
– Identify the geometry
– Considering ligands, predict the shape
Bonding (3. VSEPR Theory)
To predict molecular shapes, we will use
VSEPR Theory
Valence Shell Electron Pair Repulsion
The VSEPR model assumes that electron-electron repulsion
determines the arrangement of valence electrons around each
inner atom.
Bonding (3. VSEPR Theory)
VSEPR Theory
Electron group: set of electrons that occupies space
around an atom
LPs (2 e-), single/double/triple bonds (2/4/6 e-), lone
electrons (1 e-)
Ligand: atom bonded to inner atom
Geometry: 3D arrangement of valence e- groups
Molecular shape: how ligands, not all e- groups, are arranged
in space
Bonding (3. VSEPR Theory)
VSEPR Geometry (before shapes)
180o
109.5o
120o
Linear (2 e- gr) Trigonal Planar (3 e- gr) Tetrahedral (4 e- gr)
90o
120o
Trigonal bipyramidal (5 e- gr)
90o
Octahedral (6 e- gr)
Bonding (3. VSEPR Theory)
VSEPR Shapes: 2 e- groups
X–A–X
Linear
AX2
Example: CO2
O=C=O
AX2
Linear
Bonding (3. VSEPR Theory)
VSEPR Shapes: 3 e- groups
Remember, lines to lone pairs are NOT bonds.
This is just to show how geometry is arranged in 3D space.
Bonding (3. VSEPR Theory)
VSEPR Shapes: 3 e- groups
Example: AlCl3
AX3
Trigonal Planar
Bonding (3. VSEPR Theory)
VSEPR Shapes: 4 e- groups
Tetrahedral
AX4
Trigonal pyramidal
AX3E
Bent
AX2E2
Bonding (3. VSEPR Theory)
VSEPR Shapes: 4 e- groups
Example: CH4
AX4
Tetrahedral
Bonding (3. VSEPR Theory)
VSEPR Shapes: 4 e- groups
More Examples: NH3 , H2O
AX3E
Trigonal pyramidal
AX2E2
Bent
Bonding (3. VSEPR Theory)
VSEPR Shapes: 4 e- groups
More Examples: NH3 , H2O
Repulsion between the LPs and bonded e- maintains shape,
and slightly distorts geometry angles
Bonding (3. VSEPR Theory)
VSEPR Shapes: 4 e- groups
More Examples: NH3 , H2O (alternate views)
Bonding (3. VSEPR Theory)
VSEPR Shapes: 5 e- groups
Bonding (3. VSEPR Theory)
VSEPR Shapes: 5 e- groups
The LPs of 5 e- groups have specific patterns:
• 5 e- group geometry is the first where not all angles
are equivalent
• Place LPs around central atom to minimize
repulsion
• Remember ideal angles – use them to determine
where LPs go
Bonding (3. VSEPR Theory)
VSEPR Shapes: 5 e- groups
Where do the lone pairs go?
– Equatorial electron groups are 120o from one another
– Axial electron groups are 180o from one another, but 90o
from equatorial electron groups
• So: fill equatorial positions first because they’re
farther away from one another
90o
120o
Bonding (3. VSEPR Theory)
VSEPR Shapes: 5 e- groups
Bonding (3. VSEPR Theory)
VSEPR Shapes: 5 e- groups
Examples: SF4 , ClF3
Bonding (3. VSEPR Theory)
VSEPR Shapes: 5 e- groups
Another Example: XeF2
Bonding (3. VSEPR Theory)
VSEPR Shapes: 6 e- groups
Bonding (3. VSEPR Theory)
VSEPR Shapes: 6 e- groups
Where do the lone pairs go?
– All geometry positions are equivalent (90o)
So: put the LPs as far away from one another
as possible
– Start with 180o away
– When axial positions are occupied, must fill
equatorial
• 2nd equatorial LP should be 180o away from 1st!
Bonding (3. VSEPR Theory)
VSEPR Shapes: 6 e- groups
Bonding (3. VSEPR Theory)
VSEPR Shapes: 6 e- groups
Examples: SF6 , IF5
Bonding (3. VSEPR Theory)
VSEPR Shapes: 6 e- groups
Another Example: XeF4
Bonding (3. VSEPR Theory)
Remember:
• Molecules with no lone pairs have ideal bond
angles (as predicted by VSEPR)
• When molecules have lone pairs:
– Lone pairs alter bond angles (more e- repulsion from LPs
and bonding e-)
– More LPs = more repulsion = smaller angles
(AX4 predicts bond angle of 109.5o)
Bonding (3. VSEPR Theory)
Remember:
• Consider electronegative atoms: attract electrons
in bonds
– Allow ligands at the other end of the molecule to move
part
(AX4 predicts bond angle of 109.5o)
Bonding (3. VSEPR Theory)
VSEPR Summary Table
Electron Groups
Bonds
Lone Pairs
Notation
Shape
2
2
0
AX2
linear
3
3
0
AX3
trigonal planar
3
2
1
AX2E
bent
4
4
0
AX4
tetrahedral
4
3
1
AX3E
trigonal pyramidal
4
2
2
AX2E2
bent
5
5
0
AX5
trigonal bipyramidal
5
4
1
AX4E
see-saw
5
3
2
AX3E2
T-shaped
5
2
3
AX2E3
linear
Bonding (3. VSEPR Theory)
VSEPR Summary Table (continued)
Electron Groups
Bonds
Lone Pairs
Notation
Shape
6
6
0
AX6
octahedral
6
5
1
AX5E
square pyramidal
6
4
2
AX4E2
square planar
6
3
3
AX3E3
T-shaped
6
2
4
AX2E4
linear
Bonding (4. Theories of Chemical Bonding )
Review
Lewis structure – accounts for all
valence electrons
VSEPR – predicts the shape
around a central atom
Bonding (4. Theories of Chemical Bonding )
Most chemical bonds are polar, meaning that one
end is slightly negative and the other is slightly
positive,
A molecule with this type of
asymmetrical distribution of
electron density is said
to have Dipole Moment
Bonding (4. Theories of Chemical Bonding )
Dipole Moment (μ) depends on:
1.on bond polarities , ΔΧ (electronegativity
differences),
2. Molecular shape
We measure the polarity of a molecule as the dipole
moment (μ), with a directional indication (by words or
diagram)
Bonding (4. Theories of Chemical Bonding )
Dipole Moment (μ):
The SI unit for the dipole moment is the coulomb
meter (C m), but experimental values are usually
recorded in units of Debye,
1 Debye = 3.34x10-30 C m
+q
-q
r
μ = qr
Bonding (4. Theories of Chemical Bonding )
Dipole Moment (μ):
The polarity between bonded atoms (in a molecule)
depends on the difference of electronegativity
Example: Cl2
Cl
Cl
ΔΧ = 0
Thus μ = 0
Bonding (4. Theories of Chemical Bonding )
Dipole Moment (μ):
The polarity between bonded atoms (in a molecule)
depends on the difference of electronegativity
Another Example: HCl
ΔΧ = 1.9
μ = 1.07 D
Bonding (4. Theories of Chemical Bonding )
Dipole Moment (μ):
The polarity between bonded atoms (in a molecule)
depends on the difference of electronegativity
Another Example: CO2
ΔΧ = 1.0
ΔΧ = 1.0
μ =0, because identical polar bonds point in opposite directions,
Olmsted, Chemistry, 3e, Canadian Edition
Bonding (4. Theories of Chemical Bonding )
Dipole Moment (μ):
The polarity between bonded atoms (in a
molecule) depends on the molecular shape &
ΔΧ
Example : H2O
Bonding (4. Theories of Chemical Bonding )
Dipole Moment (μ):
The polarity between bonded atoms (in a molecule)
depends on the difference of electronegativity
Another Example: CH4
μ=0
Bonding (4. Theories of Chemical Bonding )
The polarity between bonded atoms (in a
molecule) depends on the molecular shape &
ΔΧ
Another Example:
Bonding (4. Theories of Chemical Bonding )
Review :
Olmsted, Chemistry, 3e, Canadian Edition
Bonding (4. Theories of Chemical Bonding )
VSEPR and Hybrid Orbital sites used today:
http://www.chemtube3d.com/A%20Level%20orbitals-all.htm
http://winter.group.shef.ac.uk/orbitron/MOs/N2/2s2ssigma/index.html
Go play!
Bonding (5. Orbital Overlap Model)
So far we have seen:
Atomic Theory → shapes of atomic orbitals
VSEPR → shapes of molecules
But how are atomic orbitals combined to make bonds
and give these molecular shapes?
1. Orbital Overlap Model
2. Hybrid Orbital Theory
3. Molecular Orbital Theory
Bonding (5. Orbital Overlap Model)
1. Orbital Overlap Model
Example, Consider HF:
H
1s1
F
[He] 2s2 2p5
1s ____
2p ____ ____ ____
2s ____
Bonding (5. Orbital Overlap Model)
1. Orbital Overlap Model
Example, Consider HF:
1s
2p
Bonding (5. Orbital Overlap Model)
1. Orbital Overlap Model
Another Example, Consider F2:
+
→
F2
Bonding (5. Orbital Overlap Model)
1. Orbital Overlap Model
Another Example, Consider H2S:
S
[Ne] 3s2 3p4
3p ____ ____ ____
3s ____
The H-atom 1s electrons will overlap with
these half-filled orbitals
Bonding (5. Orbital Overlap Model)
1. Orbital Overlap Model
Another Example, Consider H2S:
Bonding (5. Orbital Overlap Model)
1. Orbital Overlap Model
Another Example, Consider H2S:
Orbital overlap model → 90o
Bonding (5. Orbital Overlap Model)
1. Orbital Overlap Model
Another Example, Consider H2S:
Orbital overlap model → 90o
VSEPR → 109.5o (4 electron groups)
Actual = 92.1o
we need a better theory…
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 4:
Example: methane (CH4)
VSEPR → 109.5o
Orbital overlap model → 90o
When theory and experiment disagree, the theory
must be modified.
Steric Number = 4
(4 charge clouds)
Hybrid Orbital Theory
The steric number is the number of atoms bonded to a central atom of a molecule
plus the number of lone pairs attached to the central atom.
Bonding (5. Hybrid Orbital Theory)
In CH4, those charge clouds are
bonds between C and H
A ground state carbon atom has e- config.:
____
____
____ ____ ____
1s
2s
2p
Only 2 unpaired e- looks like only 2 bonds will form;
CH2 would be predicted!
Bonding (5. Hybrid Orbital Theory)
In CH4, those charge clouds are
bonds between C and H
An excited state carbon promotes a 2s e- to the 2p:
____
____
1s
2s
____ ____ ____
2p
Bonding (6. Hybrid Orbital Theory)
In CH4, those charge clouds are
bonds between C and H
An excited state carbon promotes a 2s e- to the 2p:
____
____
1s
2s
____ ____ ____
2p
Now 4 unpaired e- can allow 4 bonds to form;
CH4 would be predicted!
Bonding (6. Hybrid Orbital Theory)
How CH4 is made using atomic orbitals:
2p
2s
But this shape is wrong (expecting tetrahedral)
Bonding (6. Hybrid Orbital Theory)
How CH4 is made using atomic orbitals:
2p
2s
Tetrahedral
coordination
But this shape is wrong (expecting tetrahedral)
Bonding (6. Hybrid Orbital Theory)
Steric Number 4 (ex. CH4): The expected shape is
achieved by hybridization
x4
Bonding (6. Hybrid Orbital Theory)
Steric Number 4 (ex. CH4): The expected shape is
achieved by hybridization
x4
x4 sp3
(C valence)
The smaller lobe is omitted in subsequent diagrams for clarity
Bonding (6. Hybrid Orbital Theory)
Steric Number 4 (ex. CH4): The expected shape is
achieved by hybridization
x4
4 e- now shared
in hybrid orbital
Bonding (6. Hybrid Orbital Theory)
Steric Number 4 (ex. CH4): Then bonds are formed
to other atoms using those hybrid orbitals
H
H
H
C
H
C
H
H
H
H
Bonding (6. Hybrid Orbital Theory)
Also Steric Number 4:
N and O also form sp3 hybrid orbitals: N
N
1s2 2s2 2p3
2p ___ ___ ___
2s ___
(5 valence e-)
2 sp3 ___ ___ ___ ___
1 set of paired e(lone pair)
3 sp3 orbitals
used for bonds
Bonding (6. Hybrid Orbital Theory)
Also Steric Number 4:
N and O also form sp3 hybrid orbitals: N
Example: NH3
··
N
Bonding (6. Hybrid Orbital Theory)
Also Steric Number 4:
N and O also form sp3 hybrid orbitals: O
O
1s2 2s2 2p4
2p ___ ___ ___
2s ___
(6 valence e-)
2 sp3 ___ ___ ___ ___
2 sets of paired e(lone pairs)
2 sp3 orbitals
used for bonds
Bonding (6. Hybrid Orbital Theory)
Also Steric Number 4:
N and O also form sp3 hybrid orbitals: O
Example: H2O
··
··
O
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 5:
Example: phosphorous pentachloride (PCl5)
Steric Number = 5
(5 charge clouds)
The steric number is the number of atoms bonded to a central atom of
a molecule plus the number of lone pairs attached to the central atom.
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 5:
Example: phosphorous pentachloride (PCl5)
In third row we have 3d orbitals that can be involved in hybridisation
P
(1s2 2s2 2p6 3s2 3p3)
(5 valence e-)
3d ___ ___ ___ ___ ___
3p ___ ___ ___
3s ___
___
3s
___ ___ ___ ___ ___ ___ ___ ___
3p
3d
sp3d hybridization
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 5:
Example: phosphorous pentachloride (PCl5)
___
3s
___ ___ ___ ___ ___ ___ ___ ___
3p
3d
sp3d hybridization
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 6:
Example: sulfur hexafluoride (SF6)
Steric Number = 6
(6 charge clouds)
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 6:
Example: sulfur hexafluoride (SF6)
S
(1s2 2s2 2p6 3s2 3p4)
(6 valence e-)
3d ___ ___ ___ ___ ___
3p ___ ___ ___
3s ___
___
3s
___ ___ ___ ___ ___ ___ ___ ___
3p
3d
sp3d2 hybridization
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 6:
Ex: sulfur hexafluoride (SF6)
___
3s
___ ___ ___ ___ ___ ___ ___ ___
3p
3d
sp3d2 hybridization
Bonding (6. Hybrid Orbital Theory)
2p
sp2
3x sp2
hybrid
orbitals
Leftover p-orbital!
Steric Number 3
1x p orbital
(for double
bond)
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 3:
Example C2H4: (each carbon atom has 3x sp2 + 1x p)
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 3:
Example C2H4: (each carbon atom has 3x sp2 + 1x p)
C
C
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 3: Example C2H4
Direct orbital overlap = a sigma (σ) bond
(“end-to-end” overlap)
C
C
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 3: Example C2H4
Direct orbital overlap = a sigma (σ) bond
(“end-to-end” overlap)
H (1s)
C
H (1s)
C
H (1s)
H (1s)
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 3: Example C2H4
Direct orbital overlap = a sigma (σ) bond
Indirect orbital overlap = a pi (π) bond
(p)
(p)
H (1s)
(sp2)
(sp2)
(sp2)
C
H (1s)
(sp2)
(sp2)
(sp2)
C
H (1s)
H (1s)
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 3: Example C2H4
Double bond: having both a σ and a π bond between
the same atoms
(2p – 2p)
C sp2
sp2 Csp2
1s
Bonding (6. Hybrid Orbital Theory)
2p
sp
3x sp2
hybrid
orbitals
Leftover p-orbital!
Steric Number 3
1x p orbital
(for double
bond)
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 2:
Example C2H2 (HC≡CH): (each carbon atom)
2p
sp
2x sp
hybrid
orbitals
Leftover p-orbitals!
Steric Number 3
2x p orbital
(for double
bond)
Bonding (6. Hybrid Orbital Theory)
Consider Steric Number 2: Example C2H2
Triple bond: having a σ and two π bonds between the
same atoms
π
(2p – 2p)
H
σ
(1s – sp)
C
π
(2p – 2p)
σ
(sp – sp)
C
σ
(sp – 1s)
H
Bonding (6. Hybrid Orbital Theory)
Summary of hybrid orbitals
Charge Clouds
(Steric Number)
2
3
4
5
6
Hybridization
sp
sp2
sp3
sp3d
sp3d2
Bonding (6. Hybrid Orbital Theory)
So far we have seen:
• Orbital Overlap Model: localized overlap of atomic
valence shell orbitals
• Hybrid Orbital Theory: localized overlap of
hybridized and atomic orbitals
Now we will add:
• Molecular Orbital (MO) Theory: orbitals
delocalized over entire molecule
Bonding (7. Molecular Orbital Theory)
Molecular Orbital Theory (MO)
We were previously considering atomic orbitals
(coordinated between atoms only)
Now we will consider the orbitals of molecules
atomic orbitals : atom
MO : molecules
Bonding (7. Molecular Orbital Theory)
Remember in Chapter 4 & 5 we talked about
wave particle duality?
Bonding (7. Molecular Orbital Theory)
Constructive and Destructive overlap in Is
orbitals
Bonding (7. Molecular Orbital Theory)
electron density
When two orbitals have a positive e- density
overlap, they form a bonding molecular orbital
+
-
nuclei
Bonding (7. Molecular Orbital Theory)
electron density
The opposite is also possible: negative overlap
forms an antibonding molecular orbital
a node
+
-
nuclei
A “node” is a region of 0 electron density
Bonding (7. Molecular Orbital Theory)
Constructive and Destructive overlap in Is
orbitals
Bonding (7. Molecular Orbital Theory)
Bonding results in stability
Bonding (7. Molecular Orbital Theory)
Drawing MO diagrams:
1. Start with the isolate atoms
Energy
Example: H2
1s
1s
H
H
Bonding (7. Molecular Orbital Theory)
Drawing MO diagrams:
2. Build the MO’s
s*1s
Energy
Example: H2
H
s1s H
Bonding (7. Molecular Orbital Theory)
Drawing MO diagrams:
3. Put e- in the MO’s
s*1s
Energy
Example: H2
H
s1s H
Bonding (7. Molecular Orbital Theory)
s*1s
Energy
Drawing MO diagrams:
Example: H2
s1s H
H
H2
Bonding (7. Molecular Orbital Theory)
Will the molecule hang together or fall
apart?
Calculate the Bond Order:
Bonding (7. Molecular Orbital Theory)
σ1s
σ*1s
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: Another example: H2-
Energy
s*1s
s1s H
H
H2-
Bonding (7. Molecular Orbital Theory)
σ1s
σ*1s
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: Another example: H22-
s*1s
BO = 0
does not exist
s1s
H2-2
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: Another example: He2
s*1s
He
He
BO = 0
does not exist
s1s
He2
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: Another example: Li2
Li: 1s1 2s1
s*2s
2s
BO = 1
s2s
2s
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: Another example: Be2
Be: 1s1 2s2
s*2s
2s
BO = 0
(does not exist)
s2s
2s
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: Not just restricted to sorbitals – larger atoms have p-orbitals
2p
2p
s*2s
2s
2s
s2s
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: p-orbitals
2p
2p
s*2s
2s
2s
s2s
Bonding (7. Molecular Orbital Theory)
Remember p-orbitals have two possible bondtypes: σ AND π
σ-bond
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: p-orbitals
s*2p
2p
2p
s2p
Bonding (7. Molecular Orbital Theory)
Remember p-orbitals have two possible bondtypes: σ AND π
π-bonds
σ-bond
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: p-orbitals
(Note the order in bonding vs antibonding levels)
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: 2s and 2p valence
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: Example: O2 (BO = 2)
Each O has
2s2 2p4 valence
(6 valence
e- each)
BO = (8-4)/2 = 2
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: Example: F2 (BO = 1)
Each F has
2s2 2p5 valence
(7 valence
e- each)
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: Example: Ne2 (BO = 0)
Each Ne has
2s2 2p6 valence
(8 valence e-)
Bonding (7. Molecular Orbital Theory)
MO diagrams and BO: Example: N2 (BO = 3)
Each N has
2s2 2p3 (5 valence e-)
MO diagram for
molecules that
includes Nitrogen or
lighter atoms
Bonding (7. Molecular Orbital Theory)
This switch in orbital ordering occurs because of a phenomenon
called s-p mixing. The energy of the 2s and 2p orbitals in light
atoms (Li, Be, B, C, N) are close enough to each other that the
σ orbital formed from the 2s orbitals can combine with the
σ orbital formed from the 2p orbitals. This makes the σ from the
2s orbitals more stable, and the σ from the 2p orbitals less
stable. Similarly, the antibonding orbitals also undergo s-p
mixing, with the σ* from the 2s becoming more stable and the
σ* from the 2p becoming less stable.
Bonding (7. Molecular Orbital Theory)
Molecular Orbitals can explain physical properties:
• Bond energies
• Bond lengths
• Magnetic properties
Bonding (7. Molecular Orbital Theory)
Example: compare N2 to N2+
Bonding (7. Molecular Orbital Theory)
Example: compare N2 to N2+:
Bond Order
Bond Energy (kJ/mol)
Bond Length (nm)
N2
3
945
0.110
N2+
2.5
841
0.112
Bonding (7. Molecular Orbital Theory)
Another Example: compare O2 to O2+
Bonding (7. Molecular Orbital Theory)
Another example: compare O2 to O2+:
Bond Order
Bond Energy (kJ/mol)
Bond Length (nm)
O2
2
498
0.121
O2+
2.5
623
0.112
Bonding (7. Molecular Orbital Theory)
Another example: compare O2 to O2- : O2 (BO = 2)
Bonding (7. Molecular Orbital Theory)
Other Examples:
Molecule
BO
1
Bond Energy
(kJ/mol)
154
Bond Length
(nm)
0.144
F2
B2
O2
C2
N2
1
2
2
3
290
495
620
942
0.160
0.121
0.131
0.110
Bonding (7. Molecular Orbital Theory)
When drawing a Lewis Structure of NO, can
we decide between these two resonances?
. ..
: N = O:
FCN = 5 – 5 = 0
FCO = 6 – 6 = 0
or
.. .
:N = O
FCN = 5 – 6 = -1
FCO = 6 – 5 = +1
(seems more likely, but can use MO Theory to confirm)
Smaller change
is easier to reach
Single non-bonding
Π*
Bigger Lobs on Nitrogen in π*
Bonding (7. Molecular Orbital Theory)
Magnetism
• Paramagnetic: unpaired electrons are
attracted to a magnetic field
• Diamagnetic: paired electrons are repelled
by a magnetic field
We could note magnetism of elements with electron
configurations, but we care about molecules.
Bonding (7. Molecular Orbital Theory)
Magnetism: Example O2 (BO = 2)
Unpaired!
(Hund’s Rule
still applies here)
HOMO stand for Highest Occupied
Molecular Orbital .
LUMO stands for Lowest Unoccupied
Molecular Orbital.
Bonding (7. Molecular Orbital Theory)
What about something more complex, with
resonance structures? Consider:
Benzene (C6H6)
Bonding (7. Molecular Orbital Theory)
Consider: Benzene (C6H6)
Resonant forms:
Kekulé structures
=
MO
=
Bonding (7. Molecular Orbital Theory)
Consider: Benzene (C6H6)
Parallel p-orbitals all share electron density in
resonance (a.k.a. delocalized electrons)
Bonding (7. Molecular Orbital Theory)
Another example: Butadiene (C4H6)
Alternating single and double bonds
is known as “conjugation”
Bonding (7. Molecular Orbital Theory)
Another example: Butadiene (C4H6)
σ-bonds
H
H
H
C
C
C
C
H
H
π-bonds
Olmsted, Chemistry, 3e, Canadian Edition
H
Bonding (7. Molecular Orbital Theory)
Conjugation may lead to coloured compounds:
LUMO
*
h
HOMO

Unconjugated System
Large energy gap
(UV absorption)
LUMO
*
HOMO

h
Conjugated System
Smaller energy gap
(VIS absorption, coloured!)
Bonding (7. Molecular Orbital Theory)
Conjugation may lead to coloured compounds:
Examples: β-carotene
Bonding (7. Molecular Orbital Theory)
Conjugation may lead to coloured compounds:
Examples: chlorophyll
Bonding (7. Molecular Orbital Theory)
Conjugation may lead to coloured compounds:
Examples: xanthin
Bonding (7. Molecular Orbital Theory)
The idea of delocalized electrons can be applied to
metals as well: band theory of solids
Consider Li (the lightest metal):
• Crystal is an ordered structure of Li atoms
• Metallic solid has some attraction between atoms
(but not charged)
• Each Li atom has valence of 2s1
Building the Li crystal structure one atom at a time
increases the overlap of each atom’s 2s orbital
Bonding (7. Molecular Orbital Theory)
Band Theory of Solids, Example: Li
- the energy spacing between orbitals decreases as the number of delocalized
orbitals increases.
LUMO
HOMO
Bonding (7. Molecular Orbital Theory)
Band Theory of Solids: Metals conduct electricity
Olmsted, Chemistry, 3e, Canadian Edition
Bonding (7. Molecular Orbital Theory)
Band Theory of Solids: Metals conduct electricity
HOMO
(new)
HOMO
(new)
LUMO
(new)
Bonding (7. Molecular Orbital Theory)
Summary of MO Theory:
• MO’s are delocalized (not fixed to one atom)
• No need for resonance (MO accounts for
variations)
• Can predict Bond Order
• Correlations to bond length, bond energy,
magnetic and optical properties
Bonding
• Related to Textbook Chapters 6 & 7
• ORION homework (due date on cuLearn)
• WileyPlus assignment (due date on cuLearn)