Theories of Chemical Bonding
Exam #4 (Chapter 8) on 7-December
One 3” x 5” notecard
Chapter 8 OWL Deadline Tuesday Night
Chapter 9 OWL Deadline @ Final Exam
Final Exam (review session weekend before)
Monday, 12-December @ 8AM in IRC 3 (here)
~50 points just on Chapter 9
(about the same as the other chapters)
~100 points cumulative (Chapters 2-8)
One 81/2” x 11” note sheet allowed
Final Lab Experiment
This Week (let’s look @ WIKI Page)
Laboratory Checkout Next Week!
What’s Gotten Us Here….

Electron Configurations
◦ Energy Level Diagrams
◦ Valence Electrons

Lewis Structures
◦ Identify Central Atom, Bonding and
Lone Pairs
◦ Formal Charge

Periodic Trends
◦ Size, Electronegativity

Bond Properties
◦ Bond Energy, Length, Order, Polarity

Electron Pair Geometry
◦ Clues to the shape of the
molecule or ion

Molecular Geometry
◦ Definitive shape of the molecule
or ion


Bond Angles…
However, other than
identifying ionic and covalent
bonds, we haven’t really
described how atoms bond
together!
◦ Electrons in a covalent bond are
shared, but how?
Let’s Look at Boron Trifluoride
B has 3 valence e Fluorine has 7
valence e We know that B
follows an exception
to the octet rule
 Let’s look at the
Lewis Structure

How do these 3 bonds form?
We have a single s
and three p orbitals
on the boron atom
 We have a single s
and three p orbitals
on the fluorine atom
 How do they overlap
to share electrons?
 Hybridization!

Valence Bond and Molecular Orbital Theories

Share some principles
◦ Bonds are formed when electrons are shared between atoms
◦ The sharing of electrons, and their attraction to the two nuclei of
the atoms that are bonded lowers the total energy in the
molecule or ion
◦ Two types of bonds can form
◦ sigma (σ) – bonds which lie along the axis between to atoms in a
molecule or ion
 Atoms can rotate around sigma bonds
◦ pi (π) – bonds which lie in regions outside of the axis between
two atoms (but parallel to the axis)
 Atoms are fixed and can’t rotate
 Multiple bonding
Valence Bond Theory

Takes into account
◦ Lewis structure (# of bonds)
◦ VSEPR (overall shape of the molecule or ion)
◦ The available orbitals that can be used for
bonding
◦ Quantum Mechanics
◦ Spectroscopy results (like the colorimetry you
used in lab)
◦ Thermodynamic data

Helps decide what types of hybrids will form
Hybridization (very systematic)

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To obtain the bonding description about any
atom in a molecule:
1. Write the Lewis electron-dot formula.
2. Use VSEPR to determine the electron
arrangement about the atom.
3. From the arrangement, deduce the hybrid
orbitals.
4. Assign the valence electrons to the hybrid
orbitals one at a time, pairing only when
necessary.
5. Form bonds by overlapping singly occupied
hybrid orbitals with singly occupied orbitals of
another atom.
Some Tenets of Hybridization
The total number of orbitals you started with (all
valence) MUST equal the number of orbitals that
you end up with
 The number of hybrid orbitals must equal the
number of initial orbitals you started with (again
all valence)
 Two electrons per orbital
 The number of sigma bonds equals the number of
single bonds

◦ Multiple bonds are sigma + pi bonds

s, p, d orbitals are what we are working with
◦ Valence ONLY!!!
Let’s Start with Methane (CH4)

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Write the valence electron configurations
Carbon has 4 valence electrons
4 hydrogen atoms each have 1 valence eDraw the Lewis structure
Evaluate electron and molecular structure
◦ Shape helps us understand hybridization


Determine what orbitals will be combined
(hybridized) to create the necessary bonds
Describe the hybridization
Lone Pairs ?


Lone pairs of electrons can be placed in
hybrid orbitals too.
Ammonia
◦
◦
◦
◦
Write electron configurations (valence)
Lewis Structure
Predict electron and molecular geometry
Determine what orbitals might combine and
hybridize to describe the bonding
 Consider the best way to locate lone pairs (VSEPR)
◦ Describe the hybridization

Let’s do this on the board!
OK, Let’s Get Back to BF3
Lewis Structure
 Electron Pair and Molecular Structure
(geometry)
 Predict Hybridization

◦ Remember-we are normally talking about
hybridization about central atoms at this
point.

Describe Hybridization
3rd Period Elements (d orbitals)

Hybrids can be made from combinations
of s, p and d orbitals.
◦ Only available to elements in the 3rd period or
higher (n = 2 only has s and p orbitals
available)

Same rules apply
◦ Total number of orbitals are constant
◦ Number of hybrid orbitals equals the number
of orbitals combined to make them
◦ Some orbitals may be left unhybridized
5 electron pairs (sp3d)
6 electron pairs (sp3d2)
How is this systematic ?
Hybrid orbitals are named by using the atomic
orbitals that combined:
• one s orbital + one p orbital gives two sp
orbitals
• one s orbital + two p orbitals gives three sp2
orbitals
• one s orbital + three p orbitals gives four sp3
orbitals
• one s orbital + three p orbitals + one d orbital
gives five sp3d orbitals
• one s orbital + three p orbitals + two d orbitals
gives six sp3d2 orbitals
What about lone pairs
(on the central atom)

We’ve seen NH3, where they go in one of
the hybrid orbitals
◦ Usually this is the case, because the hybrid
orbital represents the greatest separation of
electron pairs (bonding or lone) in 3D space.

They can go in other hybrid orbitals in
other cases, for example…..
4 total pairs (2 structural, 2
lone)

SeH2
◦ Valence Electron Configuration
◦ Lewis Dot Structure
◦ Predict Electron Pair and Molecular Geometry
 Different-two lone pairs.
◦ Predict Hybridization
◦ Draw with electron pairs in the appropriate
orbitals and describe the hybridization
◦ What happens to the tetrahedral bond angles?

Let’s do this on the board.
5 total pairs (4 structural, one lone)
pi (π) bonds
sigma (σ) bonds lie along the internuclear axis
between atoms
 pi (π) bonds are parallel to the internuclear axis, but
lie separated from it in space

◦ Created from p orbitals not used in hybridization
◦ Responsible for multiple bonds

We will concern ourselves with pi bonds in carbon
containing (organic) molecules only
Double Bonds (one σ, one π)

Ethene (ethylene):
◦ C 2H 4
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3 x sp2 hybrid orbitals on each carbon
2 C-H sigma bonds on each carbon
1 C-C sigma bond
1 C-C pi bond
Double bond C-C (one sigma plus one pi)
Note that the geometry
of the sp2 hybrids
(trigonal planar) controls
geometry about each C
atom.
Triple Bonds (one σ, two π)

Acetylene
◦ C 2H 2





2 sp hybrid orbitals on each carbon
1 C-H sigma bonds on each carbon
1 C-C sigma bond
2 C-C pi bonds
Triple bond C-C (one sigma plus two pi)
Note that the geometry
of the sp hybrids (linear)
controls the
geometry about each C
atom.
What’s different about multiple bonds
(sigma + pi) vs. single (sigma) bonds?

Location of the bonds
◦ Along the internuclear axis versus parallel to it
Single bonds (sigma only) can rotate
 Multiple bonds (sigma + pi) are rigid

◦ Leads to structural isomers
◦ cis vs. trans forms

More than one multiple (sigma + pi) in a
molecule can lead to different resonance
forms
Do we have cis vesus trans 1, 2 dichloroethane?
Let’s draw on the board.
Benzene (C6H6) / Resonance

6 carbons, each with 3 sp2 hybrid orbitals
◦ Each results in 3 sigma bonds
◦ 1 sigma bond to H, 2 sigma bonds to C
◦ One pi bond on each carbon atom
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
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sp2 hybridization still controls geometry on each
carbon atom
pi bonds make the molecule “flat”
Two resonance forms describe possible location
of pi bonds
◦ How does this influence bond order, bond length,
bond energy?

VERY stable-electrons are distributed
throughout the molecule.
Molecular Orbital (MO) Theory
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Considers orbitals to be more spread out
(delocalized) than Valence Bond Theory
Spectroscopy results are more supportive of MO
theory
Is better at describing why certain molecules do
not form (why bonding does not happen in some
cases)
Can be used in a complimentary fashion with
Valence Bond Theory to describe what is
happening in molecules.
We still have sigma (σ) and pi (π) bonding.
◦ We add non-bonding and antibonding variables
How to conceptualize MO theory?
Instead of thinking of how hybrids form
on a central atom….
 Think about how orbitals on two
separate atoms in a bond can come
together and form bonds in a molecule

◦ Combining orbitals (MO) instead of
overlapping orbitals (VBT)
What is some of this terminology?
What do the symbols mean?
Final Exam & Reminders
OWL Deadline (Chapter 9) is the start of the
final exam
 Review Sessions (IRC 3-here)

◦ Sunday 10-11:30 AM
◦ Sunday 12:30-2 PM

Final Exam (150 points)
◦
◦
◦
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45 points Chapter 9 (short answer)
128 points Chapters 2-8 (multiple choice)
Pen/Pencils, Calculator, Ruler, 81/2” x 11” note sheet
I will provide thermodynamic and other constants
 Planck’s constant, speed of light, bond enthalpies, etc.
Some important principles of MO
Theory


The number of molecular orbitals formed
equals the number of atomic orbitals that
have combined!
Electrons are spread out over the entire
molecule
◦ As opposed to being more localized in specific
bonds as we see with VBT

The total energy of the molecular orbitals
equals the sum of the energy of the atomic
orbitals
◦ It is just that some orbitals don’t contain
electrons (antibonding)
What about He2?
H2 forms, He2 does not. Bond
Order?

Bond Order =
◦ ½ (# bonding electrons - # non-bonding electrons)

What is the bond order for H2 ?

What is the bond order for He2 ?

Bond orders less than zero suggest a molecule
would not ordinarily exist.
◦ A bond order between 0 and 1 suggests that the
molecule would also not exist or would be unstable
Other homonuclear diatomic molecules (through F2)
Nitrogen (N2)
Bond Order ?
Diamagnetic or
Paramagnetic ?
Oxygen (O2)
Bond Order ?
Diamagnetic or
Paramagnetic ?
Heteronuclear Diatomic Molecules

The total energy of the orbitals created is
still the same as the total of the atomic
orbitals, but the energy of the original
“incoming” atomic orbitals are different.

Same rules apply in terms of # of orbitals,
magnetic properties, bond order, etc.
MO vs.VB Theories…..