June 9, 2009 – Class 35 and 36 Overview
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11.3 Hybridization of Atomic Orbitals
– Hybrid orbitals properties including energy and
shape, types (sp, sp2, sp3, sp3d, sp3d2) found in
common molecules, hybrid orbitals and VSEPR
theory.
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11.4 Multiple Covalent Bonds
– Bonding using hybridization, sigma and pi
bonds.
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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Example: Why is O2 paramagnetic?
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VSEPR theory allows shape predictions
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Neither yield quantitative information about bond
lengths or energies
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
Introduction to Valence Bond Theory
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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!
Hybridization of Atomic Orbitals
In order to account for many molecular shapes, atomic orbitals
must be hybridized.
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Hybridization: refers to combining pure atomic orbitals to
generate hybrid orbitals in the valence bond approach to
covalent bonding.
Hybrid orbital: is one of a set of identical orbitals
reformulated from pure atomic orbitals and used to describe
certain covalent bonds.
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Example: Tetrahedral carbon
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.
Hybridization of Atomic Orbitals
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sp3 hybrid orbital: one of the 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.
Hybridization of Atomic Orbitals
Hybridization of Atomic Orbitals
Bonding and structure of tetrahedral
methane (CH4) – an sp3 hybridized
molecule.
Hybridization of Atomic Orbitals
Bonding and structure of trigonal
pyramidal methane (NH3) – an sp3
hybridized molecule.
Note that hybrid orbitals can
accommodate lone pair electrons as
well as bonding pairs.
Hybridization of Atomic Orbitals
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sp2 hybrid orbital: one of 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.
Hybridization of Atomic Orbitals
Hybridization of Atomic Orbitals
Bonding and structure of trigonal
planar BF3 – an sp2 hybridized
molecule.
Hybridization of Atomic Orbitals
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sp hybrid orbital: one of the pair of 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.
Hybridization of Atomic Orbitals
Hybridization of Atomic Orbitals
Bonding and structure of linear
BeCl2 – an sp hybridized
molecule.
Hybridization of Atomic Orbitals
Hybridization of Atomic Orbitals
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sp3d hybrid orbital: one of the 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.
Hybridization of Atomic Orbitals
Bonding and structure of trigonal bipyramidal
PCl5 – an sp3d hybridized molecule.
This hybridization scheme also accounts for
the shapes of seesaw, t-shaped and some
linear molecules.
Hybridization of Atomic Orbitals
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sp3d2 hybrid orbital: one of 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.
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.
Hybridization of Atomic Orbitals
Hybridization of Atomic Orbitals
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.
We can choose the likely hybridization scheme for a central
atom in a structure in the valence-bond method by:
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.
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
Problem: Describe the molecular geometry and
propose a plausible hybridization scheme for the
central atom in the ion:
(a) Cl2F+
(b) BrF4+
Multiple Covalent Bonds
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sigma (s) 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.
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pi (p) 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.
Multiple Covalent Bonds
C2H6 - Molecule should be tetrahedral (sp3 hybridized)
about each C atom.
σ bonding; end-to-end overlap of an sp3
hybridized orbital from each carbon
Multiple Covalent Bonds
C2H6 - Molecule should be tetrahedral (sp3 hybridized)
about each C atom.
C2H4 - Molecule should be trigonal planar (sp2
hybridized) about each C atom.
Multiple Covalent Bonds – C2H4 (ethylene)
Molecular shape is
determined by the orbitals
forming the s-bonds (sframework).
Rotation about the double
bond is severely restricted,
and the double bond is rigid.
Twisting one of the –CH2
groups out of plane would
reduce the amount of porbital overlap and weaken
the p bond.
Multiple Covalent Bonds – C2H4 (ethylene)
Multiple Covalent Bonds
C2H6 - Molecule should be tetrahedral (sp3 hybridized)
about each C atom.
C2H4 - Molecule should be trigonal planar (sp2
hybridized) about each C atom.
C2H2 - Molecule should be linear (sp hybridized) about
each C atom.
Multiple Covalent Bonds – C2H2 (acetylene)
Multiple Covalent Bonds – C2H2 (acetylene)
Multiple Covalent Bonds
Multiple Covalent Bonds
Problem: Describe the types of bonds and orbitals
present for:
(a) acetone (CH3)2CO
(b) HCN
(c) CO2
Problem: Hydrazine, N2H4, and carbon disulfide, CS2, form
a cyclic molecule with the following Lewis structure.
How do shape and hybridization about C and N change
when hydrazine and CS2 form this product?