1. Molecular Geometry and Bonding Theories
Chapter 9
Chapter 9 1 1 1 1

2. Homework:
9.10, 9.12, 9.26, 9.30, 9.40, 9.44, 9.52, 9.60
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3. Molecular Shapes Molecular Shapes are determined by:
Bond Distance – Distance between the nuclei of two bonded atoms along a straight line.
Bond Angle – The angle between any two bonds containing a common atom.
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4. Molecular Shapes
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5. Molecular Shapes Valence Shell Electron Pair Repulsion Theory (VSEPR)
- VSEPR theory is based on the idea that electrostatic repulsion of the electrons are reduced to a minimum when the various regions of high electron density assume positions as far apart as possible.
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6. The VSEPR Model Predicting Molecular Geometries
draw the Lewis structure
count the total number of bonding regions and lone pairs around the central atom
arrange the bonding regions and lone pairs in one of the standard geometries to minimize e--e- repulsion
multiple bounds count as a single bonding region
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7. The VSEPR Model Predicting Molecular Geometries
draw the Lewis structure
count the total number of bonding regions and lone pairs around the central atom
arrange the bonding regions and lone pairs in one of the standard geometries to minimize e--e- repulsion
multiple bounds count as a single bonding region
Chapter 9

8. Common configurations for saturated molecules.
Molecular Shapes
Valence Shell Electron Pair Repulsion Theory (VSEPR)
Common configurations for saturated molecules.
Chapter 9

9. Common configurations for saturated molecules.
Molecular Shapes
Valence Shell Electron Pair Repulsion Theory (VSEPR)
Common configurations for saturated molecules.
Regions of Density
Shape
Bond Angle
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10. Common configurations for saturated molecules.
Molecular Shapes
Valence Shell Electron Pair Repulsion Theory (VSEPR)
Common configurations for saturated molecules.
Regions of Density
Shape
Bond Angle
2 (AX2)
Linear
180o
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11. Common configurations for saturated molecules.
Molecular Shapes
Valence Shell Electron Pair Repulsion Theory (VSEPR)
Common configurations for saturated molecules.
Regions of Density
Shape
Bond Angle
2 (AX2)
Linear
180o
3 (AX3)
Trigonal Planar
120o
Chapter 9

12. Common configurations for saturated molecules.
Molecular Shapes
Valence Shell Electron Pair Repulsion Theory (VSEPR)
Common configurations for saturated molecules.
Regions of Density
Shape
Bond Angle
2 (AX2)
Linear
180o
3 (AX3)
Trigonal Planar
120o
4 (AX4)
Tetrahedral
109.5o
Chapter 9

13. Common configurations for saturated molecules.
Molecular Shapes
Valence Shell Electron Pair Repulsion Theory (VSEPR)
Common configurations for saturated molecules.
Regions of Density
Shape
Bond Angle
2 (AX2)
Linear
180o
3 (AX3)
Trigonal Planar
120o
4 (AX4)
Tetrahedral
109.5o
5 (AX5)
Trigonal Bipyramidal
90o / 120o
Chapter 9

14. Common configurations for saturated molecules.
Molecular Shapes
Valence Shell Electron Pair Repulsion Theory (VSEPR)
Common configurations for saturated molecules.
Regions of Density
Shape
Bond Angle
2 (AX2)
Linear
180o
3 (AX3)
Trigonal Planar
120o
4 (AX4)
Tetrahedral
109.5o
5 (AX5)
Trigonal Bipyramidal
90o / 120o
6 (AX6)
Octahedral
90o
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15. The VSEPR Model
Remember, you can’t “see” lone-pairs but they do take-up space.
The names of shapes of these unsaturated structures reflects what is actually seen.
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16. The VSEPR Model
Predicting Molecular Geometries
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17. The VSEPR Model
Predicting Molecular Geometries
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18. The VSEPR Model
Molecules with Expanded Valence Shells
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19. The VSEPR Model
Molecules with Expanded Valence Shells
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20. The VSEPR Model Factors Effecting Bond Angles
The bond angle decreases as the number of lone pairs on the central atom increases.
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21. The VSEPR Model Factors Effecting Bond Angles
The bond angle decreases as the number of lone pairs on the central atom increases.
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22. The VSEPR Model Factors Effecting Bond Angles
The bond angle decreases as the number of lone pairs on the central atom increases.
Lone pairs take-up more space, therefore they push-back the chemical bonds.
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23. The VSEPR Model
To minimize e--e- repulsion, lone pairs are always placed in equatorial positions.
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24. Polarity of Molecules
To determine if a molecule is polar, you need to know two things:
- polarity of the bonds in a molecule
- how the bonds are arranged
A molecule is considered polar if its center of negative and positive charge do not coincide.
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25. Polarity of Molecules
To determine if a molecule is polar, you need to know two things:
- polarity of the bonds in a molecule
- how the bonds are arranged
A molecule is considered polar if its center of negative and positive charge do not coincide.
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26. Polarity of Molecules
To determine if a molecule is polar, you need to know two things:
- polarity of the bonds in a molecule
- how the bonds are arranged
A molecule is considered polar if its center of negative and positive charge do not coincide.
Polar molecules have a dipole (a vector quantity)
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27. Polarity of Molecules
To determine if a molecule is polar, you need to know two things:
- polarity of the bonds in a molecule
- how the bonds are arranged
A molecule is considered polar if its center of negative and positive charge do not coincide.
Polar molecules have a dipole (a vector quantity)
If these dipoles act equally and in opposition to each other, the dipoles cancel-out and the molecule is considered nonpolar.
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28. Polarity of Molecules Dipole Moments of Polyatomic Molecules Example:
CO2, each C-O dipole is canceled because the molecule is linear.
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29. Polarity of Molecules Dipole Moments of Polyatomic Molecules Example:
H2O, the H-O dipoles do not cancel because the molecule is bent.
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30. Polarity of Molecules Dipole Moments of Polyatomic Molecules
Two simple rules to help determine molecular polarity (most of the time)
If there are lone pairs on the central atom – the molecule is polar.
If there are more than one type of bonds on the central atom – the molecule is polar.
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31. Polarity of Molecules
Dipole Moments of Polyatomic Molecules
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32. Covalent Bonding and Orbital Overlap
Lewis structures and VSEPR do not explain how a bond forms.
VSEPR predicts the shape of a molecule, but it does not explain how the molecule is put together.
One method to explain bonding would be Valence Bond Theory:
Bonds form when atomic orbitals on atoms overlap.
Two electrons are shared by the orbital overlap.
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33. Covalent Bonding and Orbital Overlap
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34. Covalent Bonding and Orbital Overlap
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35. Hybrid Orbitals sp Hybrid Orbitals Consider BeF2
- Be has a 1s22s2 electron configuration.
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).
BUT the geometry is still not explained.
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36. Hybrid Orbitals sp Hybrid Orbitals
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 that come from an s-orbital and a p-orbital are called sp hybrid orbitals.
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37. Hybrid Orbitals 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.
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38. Hybrid Orbitals Hybrid Orbitals
- Other hybrids can be formed by mixing additional p and/or d orbitals
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39. Hybrid Orbitals
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40. Hybrid Orbitals
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41. Hybrid Orbitals Summary To assign hybridization:
draw a Lewis structure
assign the electron pair geometry using VSEPR theory
from the electron pair geometry, determine the hybridization
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42. Hybrid Orbitals Summary To assign hybridization:
draw a Lewis structure
assign the electron pair geometry using VSEPR theory
from the electron pair geometry, determine the hybridization
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43. Multiple Bonds
-Bonds - electron density lies on the axis between the nuclei.
-Bonds - electron density lies above and below the plane of the nuclei.
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.
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44. Multiple Bonds
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45. Multiple Bonds
Ethylene, C2H4
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46. Multiple Bonds
Ethylene, C2H4
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47. Multiple Bonds
Ethylene, C2H4
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48. Multiple Bonds
Acetylene, C2H2
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49. Molecular Orbitals
Some aspects of bonding are not explained by Lewis structures, VSEPR theory and hybridization.
For these molecules, we use Molecular Orbital Theory (MO theory).
Just as electrons in atoms are found in atomic orbitals, electrons in molecules are found in molecular orbitals.
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50. Molecular Orbitals 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).
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51. Molecular Orbitals
The Hydrogen Molecule
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52. Molecular Orbitals
The Hydrogen Molecule
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53. Molecular Orbitals
The Helium Molecule
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54. Molecular Orbitals The Hydrogen Molecule
MO diagram shows the energies and electrons in an orbital.
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.
H2 has two bonding electrons.
He2 has two bonding electrons and two antibonding electrons.
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55. Molecular Orbitals Bond Order
Bond Order = ½(bonding electrons - antibonding electrons).
Bond order = 1 for single bond.
Bond order = 2 for double bond.
Bond order = 3 for triple bond.
Bond order for H2 = ½(bonding electrons - antibonding electrons)
= ½(2 - 0) = 1.
Therefore, H2 has a single bond.
Bond order for He2 = ½(bonding electrons - antibonding electrons) = ½(2 - 2) = 0.
Therefore He2 is not a stable molecule.
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56. Second-Row Diatomic Molecules
AO’s combine according to the following rules
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;
Pauli: each MO has at most two electrons
Hund: for degenerate orbitals, each MO is first occupied singly.
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57. Second-Row Diatomic Molecules
Molecular Orbitals from 2p Atomic Orbitals
There are two ways in which two p orbitals overlap
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  orbitals
one * and two *orbitals
The relative energies of these  and  orbitals can change.
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58. Second-Row Diatomic Molecules
Molecular Orbitals from 2p Atomic Orbitals
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59. Second-Row Diatomic Molecules
Electron Configurations for B2 through Ne2
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60. Second-Row Diatomic Molecules
Electron Configurations for B2 through Ne2
As the atomic number decreases, it becomes more likely that a 2s orbital can interact with the 2p orbital.
As the 2s-2p interaction increases, the 2s MO decreases in energy and the 2p orbital increases in energy.
For B2, C2 and N2 the 2p orbital is higher in energy than the 2p.
For O2, F2 and Ne2 the 2p orbital is lower in energy than the 2p.
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61. Second-Row Diatomic Molecules
Electron Configurations for B2 through Ne2
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62. Second-Row Diatomic Molecules
Electron Configurations for B2 through Ne2
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63. Second-Row Diatomic Molecules
Electron Configurations and Molecular Properties
Two types of magnetic behavior:
Paramagnetism, unpaired electrons in molecule
Diamagnetism, no unpaired electrons in molecule
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64. Second-Row Diatomic Molecules
Electron Configurations and Molecular Properties
Chapter 9

65. Homework:
9.10, 9.12, 9.26, 9.30, 9.40, 9.44, 9.52, 9.60
Chapter 9 54 54 54 54