Transcript General Chemistry
Bonding and Molecular Structure
Chapter 10
Chapter 10 1
Covalent Bonding and Orbital Overlap
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Lewis structures and VSEPR do not explain how a
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bond forms.
VSEPR predicts the shape of a molecule, but it does
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not explain how the molecule is put together.
One method to explain bonding would be Valence Bond Theory:
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Bonds form when atomic orbitals on atoms overlap.
Two electrons are shared by the orbital overlap.
Chapter 10 2
Covalent Bonding and Orbital Overlap
Chapter 10 3
Hybrid Orbitals
sp Hybrid Orbitals
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Consider BeF 2 - Be has a 1s 2 2s 2 electron configuration.
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There is no unpaired electron available for bonding. We could promote and electron from the 2s orbital on Be to the 2p orbital to get two unpaired electrons for bonding.
The F-Be-F bond angle is 180
(VSEPR theory).
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BUT the geometry is still not explained .
Chapter 10 4
Hybrid Orbitals
sp Hybrid Orbitals
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We can solve the problem by allowing the 2s and one 2p orbital on Be to mix or form two hybrid orbitals (process called hybridization).
- The hybrid orbitals come from an s and a p orbital and is called sp hybrid orbital.
Chapter 10 5
Hybrid Orbitals
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sp Hybrid Orbitals The two sp hybrid orbitals are 180
apart.
Since only one of the Be 2p orbitals has been used in hybridization, there are two unhybridized p orbitals remaining on Be.
Chapter 10 6
Hybrid Orbitals
Hybrid Orbitals - Other hybrids can be formed by mixing additional p and/or d orbitals Chapter 10 7
Hybrid Orbitals
Chapter 10 8
Hybrid Orbitals
Chapter 10 9
Hybrid Orbitals
Summary To assign hybridization:
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draw a Lewis structure assign the electron pair geometry using VSEPR theory from the electron pair geometry, determine the hybridization Chapter 10 10
Multiple Bonds
-Bonds - electron density lies on the axis between the nuclei.
-Bonds - electron density lies above and below the plane of the nuclei.
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A double bond consists of one
-bond and one
-bond A triple bond has one
-bond and two
-bonds
-bonds come from unhybridized p orbitals.
Chapter 10 11
Multiple Bonds
Chapter 10 12
Multiple Bonds
Ethylene, C 2 H 4 Chapter 10 13
Multiple Bonds
Acetylene, C 2 H 2
Chapter 10 14
Molecular Orbitals
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Some aspects of bonding are not explained by Lewis
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structures, VSEPR theory and hybridization. For these molecules, we use Molecular Orbital Theory
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(MO theory).
Just as electrons in atoms are found in atomic orbitals, electrons in molecules are found in molecular orbitals.
Chapter 10 15
Molecular Orbitals
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The Hydrogen Molecule When two AOs overlap two MO’s form.
One molecular orbital has electron density between nuclei (
, bonding MO); One molecular orbital has little electron density between nuclei (
* , antibonding MO).
Chapter 10 16
Molecular Orbitals
The Hydrogen Molecule Chapter 10 17
Molecular Orbitals
The Hydrogen Molecule Chapter 10 18
Molecular Orbitals
The Hydrogen Molecule Chapter 10 19
Molecular Orbitals
The Hydrogen Molecule
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MO diagram shows the energies and orbitals in an MO diagram.
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The total number of electrons in all atoms are placed in the MO’s starting from lowest energy (
1s ) and ending when you run out of electrons.
Note that electrons in MO’s have opposite spins.
H 2 has two bonding electrons.
He 2 has two bonding electrons and two antibonding electrons.
Chapter 10 20
Molecular Orbitals
Bond Order Define Bond Order = ½( bonding electrons Bond order = 1 for single bond.
Bond order = 2 for double bond.
Bond order = 3 for triple bond.
antibonding electrons ).
Bond order for H 2 = ½( bonding electrons - antibonding electrons ) = ½(2 - 0) = 1.
Therefore, H 2 Bond order for He 2 has a single bond.
= ½( bonding electrons - antibonding electrons ) = ½(2 - 2) = 0. Therefore He 2 is not a stable molecule.
Chapter 10 21
Second-Row Diatomic Molecules
AO’s combine according to the following rules
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The number of MO’s equals the number of AO’s AO’s of similar energy combine As overlap increases, the energy of the MO decreases
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Pauli: each MO has at most two electrons Hund: for degenerate orbitals, each MO is first occupied singly.
Chapter 10 22
Second-Row Diatomic Molecules
Molecular Orbitals from 2p Atomic Orbitals
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There are two ways in which two p orbitals overlap
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end on so that the resulting MO has electron density on the axis between nuclei (
orbital) sideways so that the resulting MO has electron density above and below the axis between nuclei (
orbital).
The p-orbitals must give rise to 6 MO’s: one
and two
’s one
* and two
*’s The relative energies of these
and
orbitals can change.
Chapter 10 23
Second-Row Diatomic Molecules
Molecular Orbitals from 2p Atomic Orbitals Chapter 10 24
Second-Row Diatomic Molecules
Electron Configurations for B 2 through Ne 2 Chapter 10 25
Second-Row Diatomic Molecules
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Electron Configurations for B 2 For B 2 , C 2 and N 2 the
2p through Ne 2 orbital is higher in energy than the
2p .
For O 2 , F 2 than the
2p .
and Ne 2 the
2p orbital is lower in energy Chapter 10 26
Second-Row Diatomic Molecules
Electron Configurations for B 2 through Ne 2 Chapter 10 27
Second-Row Diatomic Molecules
Electron Configurations for B 2 through Ne 2 Chapter 10 28
Second-Row Diatomic Molecules
Electron Configurations and Molecular Properties Two types of magnetic behavior:
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Paramagnetism, unpaired electrons in molecule Diamagnetism, no unpaired electrons in molecule Chapter 10 29
Second-Row Diatomic Molecules
Electron Configurations and Molecular Properties Chapter 10 30
Homework
10, 12, 24, 47
Chapter 10 31