Chem 1011
Dr. L. Dawe
Chem 1011
Dr. L. Dawe
Winter 2010
January 27, 2010 – Lecture 9
10.6 Valence Bond Theory: Orbital Overlap as a
Chemical Bond
10.7 Valence Bond Theory: Hybridization of Atomic
Orbitals
– sp3 Hybridization
– sp2 Hybridization and Double Bonds
– sp Hybridization and Triple Bonds
– sp3d and sp3d2 Hybridization
– Writing Hybridization and Bond Schemes
Chem 1011
Dr. L. Dawe
Winter 2010
Lecture 9
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What A Bonding Theory Should Do
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Lewis theory is simple and structures can be
determined rapidly.
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It does not account for odd-electron species, resonance
structures or the magnetic and spectral properties of
molecules.
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VSEPR theory allows shape predictions
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Neither yield quantitative information about bond
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lengths or energies
Chem 1011
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Dr. L. Dawe
Winter 2010
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Introduction to Valence Bond Theory
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Valence-bond method: treats a covalent bond in terms
of the overlap of pure or hybridized orbitals. Electron
probability (or electron charge density) is concentrated
in the area of overlap.
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This theory tells us what a covalent bond is and correlates
molecular shapes to the interactions of atomic orbitals.
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The basic principle of valence bond theory is that a covalent
bond forms when half filled orbitals on two different atoms
(atomic orbitals) overlap. Example: H2
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Introduction to Valence Bond Theory
Lecture 9
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http://cwx.prenhall.com/petrucci/medialib/medi
a_portfolio/text_images/038_H2BondForm.MO
V
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Chem 1011
Dr. L. Dawe
Winter 2010
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Introduction to Valence Bond Theory
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Chem 1011
Dr. L. Dawe
Winter 2010
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Introduction to Valence Bond Theory
•
Localized electron model: according to valence bond theory,
core electrons and lone-pair electrons retain the same
orbital locations as in the separated atoms.
Charge density of the bonding electrons is concentrated in
regions of orbital overlap. Example: Bonding in H2S.
•
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Note: (+) and (-) signs denote phase signs, not charges!
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Introduction to Valence Bond Theory
Lecture 9
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Using the Valence-Bond Method to Describe a Molecular
Structure.
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Describe the phosphine molecule, PH3, by the valence-bond
method..
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Step 1: Draw valence shell orbital diagrams for the
separate atoms.
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Introduction to Valence Bond Theory
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Step 2: Sketch the orbitals of the central atom (P) that are
involved in the overlap.
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Recall:
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Introduction to Valence Bond Theory
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Step 3: Complete the structure by bringing together the
bonded atoms and representing the orbital overlap.
Step 4: Describe the structure.
PH3 is trigonal-pyramidal.
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Three H atoms lie in the same plane.
The P is situated at the top of the pyramid.
The three H-P-H bond angles are 90o.
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(Experimentally, they are measured to be
between 93 and 94o, vs VSEPR that
predicts slightly less than 109.5o.)
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Introduction to Valence Bond Theory
•
Note that this simple approach does not explain
bonding in methane!
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
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Hybrid orbital properties
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The number of hybrid orbitals equals the total number of
atomic orbitals that are combined.
Hybridization rationalizes experimentally determined
shape, it is not an actual physical phenomenon.
Atomic orbital energy is conserved upon hybridization.
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Example: For tetrahedral C, the p orbitals each move down
¼ of the energy difference between the s and p orbitals,
while the s orbitals move up by ¾.
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
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sp3 hybrid orbital: these are four orbitals formed by the
hybridization of one s and three p orbitals. The
angle between any two of the orbitals is the
tetrahedral angle, 109.5o.
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Think “four electron groups four hybrid orbitals”
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
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Bonding and structure of tetrahedral
methane (CH4) – an sp3 hybridized
molecule.
Note that we are only considering the
hybridization of the central atom’s
orbitals!
In the treatment in Chem 1011, we
assume the terminal atoms are
unhybridized.
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
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Bonding and structure of trigonal
pyramidal methane (NH3) – an sp3
hybridized molecule.
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Note that hybrid orbitals can
accommodate lone pair electrons as
well as bonding pairs.
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals
sp2 hybrid orbital: these are the three orbitals formed by
the hybridization of one s and two p orbitals. The
angle between any two of the orbitals is 120o.
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This hybridization scheme is common to most boron
containing compounds.
Think “three electron groups three hybrid orbitals”
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
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Bonding and structure of trigonal
planar BF3 – an sp2 hybridized
molecule.
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals
sp hybrid orbital: these are the two orbitals formed by
the hybridization of one s and one p orbital. The
angle between the two orbitals is 180o.
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This hybridization scheme is common to most beryllium
containing compounds.
Think “two electron groups two hybrid orbitals”
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
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Bonding and structure of linear
BeCl2 – an sp hybridized
molecule.
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals
Lecture 9
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http://cwx.prenhall.com/petrucci/medialib/medi
a_portfolio/text_images/056_Hybridization.MO
V
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
sp3d hybrid orbital: these are five orbitals formed by the
hybridization of one s, three p, and one d orbital.
The five orbitals are directed to the corners of a
trigonal bipyramid.
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Example: Hybridization of phosphorus, P.
Think “five electron groups five hybrid orbitals”
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
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Bonding and structure of trigonal
bipyramidal PCl5 – an sp3d hybridized
molecule.
This hybridization scheme also
accounts for the shapes of seesaw, tshaped and some linear molecules.
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals
sp3d2 hybrid orbital: these are the six orbitals formed by
the hybridization of one s, three p and two d
orbitals. The six orbitals are directed to the
corners of a regular octahedron.
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Think “six electron groups six hybrid orbitals”
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
Bonding and structure of octahdral
SF6 – an sp3d2 hybridized molecule.
This hybridization scheme also
accounts for the shapes of square
pyramidal and square planar
molecules.
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
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Note: Five
electron
groups Five sp3d
orbitals
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Note: Six
electron
groups Six sp3d2
orbitals
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybrid Orbitals and VSEPR Theory
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VSEPR method uses empirical data to give an
approximate molecular geometry, whereas the valence
bond method relates to the orbitals used in bonding
based on a given geometry.
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We can choose the likely hybridization scheme for a central
atom in a structure in the valence-bond method by:
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1.
2.
3.
writing a plausible Lewis structure for the species of interest
using VSEPR theory to predict the probable electron-group
geometry of the central atom.
selecting the hybridization scheme corresponding to the
electron-group geometry.
Chem 1011
Dr. L. Dawe
Winter 2010
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Hybrid Orbitals and VSEPR Theory
Problem: Predict the shape of the following molecules
and a hybridization scheme consistent with this
prediction
(a) SiF4
(b) XeF4
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Problem: Describe the molecular geometry and
propose a plausible hybridization scheme for the
central atom in the ion:
(a) Cl2F+
(b) BrF4+
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Chem 1011
Dr. L. Dawe
Lecture 9
Problem: Predict the shape of the following molecules and a hybridization scheme consistent with this
prediction
(a) SiF4
(b) XeF4
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Chem 1011
Dr. L. Dawe
Lecture 9
Problem: Describe the molecular geometry and propose a plausible hybridization scheme for the central
atom in the ion:
(a) Cl2F+
(b) BrF4+
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Multiple Covalent Bonds
sigma (σ) bonds: results from the end-to-end overlap of
simple or hybridized atomic orbitals along the
straight line joining the nuclei of the bonded atoms.
pi (π) bonds: results from the side-to-side overlap of p
orbitals, producing a high electron charge density
above and below the line joining the bonded
atoms.
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
Multiple Covalent Bonds
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Multiple Covalent Bonds
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C2H6 - Molecule should be tetrahedral (sp3
hybridized) about each C atom.
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σ bonding; end-to-end overlap of an
hybridized orbital from each carbon
sp3
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Multiple Covalent Bonds
Lecture 9
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C2H6 - Molecule should be tetrahedral
hybridized) about each C atom.
(sp3
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C2H4 - Molecule should be trigonal planar
(sp2 hybridized) about each C atom.
Chem 1011
Dr. L. Dawe
Winter 2010
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Multiple Covalent Bonds
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Molecular shape is
determined by the orbitals
forming the σ-bonds (σframework).
Rotation about the double
bond is severely restricted,
and the double bond is rigid.
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Twisting one of the –CH2
groups out of plane would
reduce the amount of porbital overlap and weaken
the π bond.
Chem 1011
Dr. L. Dawe
Winter 2010
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Multiple Covalent Bonds
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Multiple Covalent Bonds
Lecture 9
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C2H6 - Molecule should be tetrahedral
hybridized) about each C atom.
(sp3
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C2H4 - Molecule should be trigonal planar
(sp2 hybridized) about each C atom.
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C2H2 - Molecule should be linear (sp
hybridized) about each C atom.
Chem 1011
Dr. L. Dawe
Winter 2010
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Multiple Covalent Bonds
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Chem 1011
Dr. L. Dawe
Winter 2010
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Multiple Covalent Bonds
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15
Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Multiple Covalent Bonds
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
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Multiple Covalent Bonds
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Problem: Describe the types of bonds and orbitals
present for:
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(a) HCN
(b) CO2
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Chem 1011
Dr. L. Dawe
Winter 2010
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Looking Ahead: January 29, 2010 – Lecture 10
11.2 Solids, Liquids, and Gases: A Molecular
Comparison
– Changes between phases
11.3 Intermolecular Forces: The Forces That Hold
Condensed Phases Together
– Dispersion Force
– Dipole-Dipole Force
– Hydrogen Bonding
– Ion Dipole Force
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Chem 1011
Dr. L. Dawe
Lecture 9
Problem: Describe the types of bonds and orbitals present for:
(a) HCN
(b) CO2
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