1. Molecular Geometry and Chemical Bonding Theory

2. Bond Theory
In this chapter we will discuss the geometries of molecules in terms of their electronic structure.
We will also explore two theories of chemical bonding: valence bond theory and molecular orbital theory.
Molecular geometry is the general shape of a molecule, as determined by the relative positions of the atomic nuclei.
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3. The Valence-Shell Electron Pair Repulsion Model
The valence-shell electron pair repulsion (VSEPR) model predicts the shapes of molecules and ions by assuming that the valence shell electron pairs are arranged as far from one another as possible.
To predict the relative positions of atoms around a given atom using the VSEPR model, you first note the arrangement of the electron pairs around that central atom.
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4. Predicting Molecular Geometry
The following rules and figures will help discern electron pair arrangements.
Draw the Lewis structure
Determine how many electrons pairs are around the central atom. Count a multiple bond as one pair.
Arrange the electrons pairs are shown in Figure 10.3.
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5. Arrangement of Electron Pairs About an Atom
Linear
3 pairs
Trigonal planar
4 pairs
Tetrahedral
5 pairs
Trigonal bipyramidal
6 pairs
Octahedral
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6. Predicting Molecular Geometry
The following rules and figures will help discern electron pair arrangements.
Obtain the molecular geometry from the directions of bonding pairs, as shown in Figures 10.4 and 10.8.
(See Animations: Electron Pair Repulsion, 2 Pairs; Electron Pair Repulsion, 3 Pairs; and Electron Pair Repulsion, 4 Pairs)
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8. Fig. 10.4
Go to 10.8
Return
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9. Predicting Molecular Geometry
Two electron pairs (linear arrangement). :
You have two double bonds, or two electron groups about the carbon atom.
Thus, according to the VSEPR model, the bonds are arranged linearly, and the molecular shape of carbon dioxide is linear. Bond angle is 180o.
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10. Predicting Molecular Geometry
Three electron pairs (trigonal planar arrangement). Cl C : O
The three groups of electron pairs are arranged in a trigonal plane. Thus, the molecular shape of COCl2 is trigonal planar. Bond angle is 120o.
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11. Predicting Molecular Geometry
Three electron pairs (trigonal planar arrangement). O :
Ozone has three electron groups about the central oxygen. One group is a lone pair.
These groups have a trigonal planar arrangement.
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12. Predicting Molecular Geometry
Three electron pairs (trigonal planar arrangement). O :
Since one of the groups is a lone pair, the molecular geometry is described as bent or angular.
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13. Predicting Molecular Geometry
Three electron pairs (trigonal planar arrangement). O :
Note that the electron pair arrangement includes the lone pairs, but the molecular geometry refers to the spatial arrangement of just the atoms.
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14. Predicting Molecular Geometry
Four electron pairs (tetrahedral arrangement). :
:Cl: H N : : O H : :
:Cl C
Cl: : :
:Cl: :
Four electron pairs about the central atom lead to three different molecular geometries.
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15. Predicting Molecular Geometry
Four electron pairs (tetrahedral arrangement). :
:Cl: H N : : O H C
:Cl : :
Cl:
:Cl: :
tetrahedral
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16. Predicting Molecular Geometry
Four electron pairs (tetrahedral arrangement). :
:Cl: H N : : O H C
:Cl : :
Cl:
:Cl: :
tetrahedral
trigonal pyramid
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17. Predicting Molecular Geometry
Four electron pairs (tetrahedral arrangement). :
:Cl: : H N : C O
:Cl : : :
Cl: H
:Cl: H :
tetrahedral
trigonal pyramid
bent
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18. Predicting Molecular Geometry
Five electron pairs (trigonal bipyramidal arrangement).
: F : :
F :
: F P
This structure results in both 90o and 120o bond angles.
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19. Predicting Molecular Geometry
Other molecular geometries are possible when one or more of the electron pairs is a lone pair.
SF4
ClF3
XeF2
Let’s try their Lewis structures.
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20. Predicting Molecular Geometry
Other molecular geometries are possible when one or more of the electron pairs is a lone pair. F
ClF3
XeF2 F : S F F
see-saw
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21. Predicting Molecular Geometry
Other molecular geometries are possible when one or more of the electron pairs is a lone pair. S F : Cl F :
XeF2
see-saw
T-shape
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22. Predicting Molecular Geometry
Other molecular geometries are possible when one or more of the electron pairs is a lone pair. S F : Cl F : F : Xe : : F
see-saw
T-shape
linear
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23. Predicting Molecular Geometry
Six electron pairs (octahedral arrangement).
:F: : :F : F: : S : F: : :F
:F: :
This octahedral arrangement results in 90o bond angles.
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24. Predicting Molecular Geometry
Six electron pairs (octahedral arrangement).
IF5
XeF4
Six electron pairs also lead to other molecular geometries.
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25. Predicting Molecular Geometry
Six electron pairs (octahedral arrangement). F F F
XeF4 I F F :
square pyramid
(See Animation: Iodine Peutafluoride Structure)
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26. Predicting Molecular Geometry
Six electron pairs (octahedral arrangement). : I F : F F Xe F F :
square pyramid
square planar
Figures 10.2, 10.4, and 10.8 summarize all the possible molecular geometries.
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27. Bond Angle Examples
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28. Bond Angles and Multiple Bonds
Compare with BF3
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29. VSEPR Summary Click here for animation
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30. Dipole Moment and Molecular Geometry
The dipole moment is a measure of the degree of charge separation in a molecule.
We can view the polarity of individual bonds within a molecule as vector quantities.
Thus, molecules that are perfectly symmetric have a zero dipole moment. These molecules are considered nonpolar. (see Table 10.1) d- d+
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31. Polar Covalent Bonds
A polar covalent bond is one in which the bonding electrons spend more time near one of the two atoms involved.
When the atoms are alike, as in the H-H bond of H2 , the bonding electrons are shared equally (a nonpolar covalent bond).
When the two atoms are of different elements, the bonding electrons need not be shared equally, resulting in a “polar” bond. 2

32. Polar Covalent Bonds Fig. 5.12
(a) In the nonpolar covalent bond present, there is a symmetrical distribution of electron density. (b) In the polar covalent bond present, electron density is displaced because of its electronegativity.

33. Polarity of Covalent Bonds
In a polar covalent bond:
the more electronegative atom attracts the shared electrons more strongly and acquires a partial negative charge; indicated by d- or the head of a crossed arrow.
the less electronegative atom attracts the shared electrons less strongly and acquires a partial positive charge; indicated by d+ or the tail of a crossed arrow.
Polar Molecules

34. Polar Covalent Bonds Fig. 5.13
(a) Methane is a nonpolar tetrahedral molecule.
(b) Methyl chloride is a polar tetrahedral molecule.

35. Dipole Moment and Molecular Geometry
However, molecules that exhibit any asymmetry in the arrangement of electron pairs would have a nonzero dipole moment. These molecules are considered polar.
(See Animation: Polar Molecules) d- d+ H N : d+ d-
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36. Bond Dipoles in NF3
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37. Dipole Moment Symmetric Molecules 1.73 D Ammonia 0.44 D Trifluoroamine
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38. Go To Linus Pauling and polar covalent bonds
Dipole Moment
Glycine
Asymmetric Molecule
Go To Linus Pauling and polar covalent bonds
2.58 D
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39. Chapter 11 Valence Bond Theory
Valence bond theory is an approximate theory to explain the covalent bond from a quantum mechanical view.
According to this theory, a bond forms between two atoms when the following conditions are met.
Two atomic orbitals “overlap”
The total number of electrons in both orbitals is no more than two.
(See Figure )
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40. Hybrid Orbitals
In general, one might expect the number of bonds formed by an atom to be equal to the # of unpaired electrons.
Chlorine, for example, generally forms one bond and has one unpaired electron.
Oxygen, with two unpaired electrons, usually forms two bonds.
However, carbon, with only two unpaired electrons, generally forms four bonds. For example, methane, CH4, is well known.
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41. Hybrid Orbitals The bonding in carbon might be explained as follows:
When an electron from the 2s orbital is promoted (excited) to the vacant 2p orbital, it results in the formation of 4 unpaired electrons.
The following electronic configurations illustrate this excitation.
1s22s22p s22s12p3
2s22px12py12pz s12px12py12pz1
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42. 2p 2p 2s 2s Energy 1s 1s C atom (ground state) C atom (promoted)

43. Hybrid Orbitals
In this scenario, one bond on carbon would form using the 2s orbital while the other three bonds would use the 2p orbitals.
However, this does not explain the fact that the four bonds in CH4 appear to be identical.
Valence bond theory assumes that the four atomic orbitals in carbon combine to make four equivalent “hybrid” orbitals.
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44. Hybrid Orbitals
Hybrid orbitals are orbitals used to describe bonding that are obtained by taking combinations of atomic orbitals of an isolated atom.
In this case, a set of hybrids are constructed from one “s” orbital and three “p” orbitals, so they are called sp3 hybrid orbitals.
The four sp3 hybrid orbitals take the shape of a tetrahedron (See Figure).
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45. You can represent the hybridization of carbon in CH4 as follows.
sp3 1s
sp3 1s
Energy
The formation of two additional covalent bonds is more than enough energy to account for sp3-hybridization.
C-H bonds 1s
C atom
(excited state)
C atom
(hybridized state)
C atom
(in CH4) 2

46. Hybrid Orbitals
Note that there is a relationship between the type of hybrid orbitals and the geometric arrangement of those orbitals.
Thus, if you know the geometric arrangement, you know what hybrid orbitals to use in the bonding description.
Figure summarizes the types of hybridization and their spatial arrangements.
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47. Geometric Arrangements
Hybrid Orbitals
Hybrid Orbitals
Geometric Arrangements
Number of Orbitals
Example sp
Linear
(See Animation: sp Hydridization) 2
Be in BeF2
sp2
Trigonal planar
(See Animation: sp2 Hydridization) 3
B in BF3
sp3
Tetrahedral
(See Animation: sp3 Hydridization) 4
C in CH4
sp3d
Trigonal bipyramidal 5
P in PCl5
sp3d2
Octahedral 6
S in SF6 2

48. Hybrid Orbitals
To obtain the bonding description of any atom in a molecule, you proceed as follows:
Write the Lewis electron-dot formula for the molecule.
From the Lewis formula, use the VSEPR theory to determine the arrangement of electron pairs around the atom.
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49. Hybrid Orbitals
To obtain the bonding description of any atom in a molecule, you proceed as follows:
From the geometric arrangement of the electron pairs, obtain the hybridization type (see Table 10.2).
Assign valence electrons to the hybrid orbitals of this atom one at a time, pairing only when necessary.
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50. Hybrid Orbitals
To obtain the bonding description of any atom in a molecule, you proceed as follows:
Form bonds to this atom by overlapping singly occupied orbitals of other atoms with the singly occupied hybrid orbitals of this atom.
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51. A Problem to Consider
Describe the bonding in H2O according to valence bond theory. Assume that the molecular geometry is the same as given by the VSEPR model.
From the Lewis formula for a molecule, determine its geometry about the central atom using the VSEPR model.
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52. A Problem to Consider
Describe the bonding in H2O according to valence bond theory. Assume that the molecular geometry is the same as given by the VSEPR model.
The Lewis formula for H2O is
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53. A Problem to Consider
Describe the bonding in H2O according to valence bond theory. Assume that the molecular geometry is the same as given by the VSEPR model.
From this geometry, determine the hybrid orbitals on this atom, assigning its valence electrons to these orbitals one at a time.
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54. A Problem to Consider
Describe the bonding in H2O according to valence bond theory. Assume that the molecular geometry is the same as given by the VSEPR model.
Note that there are four pairs of electrons about the oxygen atom.
According to the VSEPR model, these are directed tetrahedrally, and from the previous table you see that you should use sp3 hybrid orbitals.
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55. A Problem to Consider
Describe the bonding in H2O according to valence bond theory. Assume that the molecular geometry is the same as given by the VSEPR model.
Each O-H bond is formed by the overlap of a 1s orbital of a hydrogen atom with one of the singly occupied sp3 hybrid orbitals of the oxygen atom.
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56. You can represent the bonding to the oxygen atom in H2O as follows:
sp3 1s
sp3 1s
Energy
lone
pairs
O-H
bonds 1s
O atom
(ground state)
O atom
(hybridized state)
O atom
(in H2O) 2

57. A Problem to Consider
Describe the bonding in XeF4 using hybrid orbitals.
From the Lewis formula for a molecule, determine its geometry about the central atom using the VSEPR model.
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58. A Problem to Consider
Describe the bonding in XeF4 using hybrid orbitals.
The Lewis formula of XeF4 is
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59. A Problem to Consider
Describe the bonding in XeF4 using hybrid orbitals.
From this geometry, determine the hybrid orbitals on this atom, assigning its valence electrons to these orbitals one at a time.
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60. A Problem to Consider
Describe the bonding in XeF4 using hybrid orbitals.
The xenon atom has four single bonds and two lone pairs. It will require six orbitals to describe the bonding.
This suggests that you use sp3d2 hybrid orbitals on xenon.
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61. A Problem to Consider
Describe the bonding in XeF4 using hybrid orbitals.
Each Xe-F bond is formed by the overlap of a xenon sp3d2 hybrid orbital with a singly occupied fluorine 2p orbital.
You can summarize this as follows:
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62. 5d 5p 5s Xe atom (ground state)2

63. Xe atom (hybridized state)
sp3d2
Xe atom (hybridized state) 2

64. 5d sp3d2 lone pairs Xe-F bonds Xe atom (in XeF4)

65. Multiple Bonding
According to valence bond theory, one hybrid orbital is needed for each bond (whether a single or multiple) and for each lone pair.
For example, consider the molecule ethene.
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66. Multiple Bonding
Each carbon atom is bonded to three other atoms and no lone pairs, which indicates the need for three hybrid orbitals.
This implies sp2 hybridization.
The third 2p orbital is left unhybridized and lies perpendicular to the plane of the trigonal sp2 hybrids.
The following slide represents the sp2 hybridization of the carbon atoms.
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67. 2p 2p sp2 2s Energy 1s 1s C atom (ground state) C atom (hybridized)
(unhybridized) 2p 2p
sp2 2s
Energy 1s 1s
C atom (ground state)
C atom (hybridized) 2

68. Multiple Bonding
To describe the multiple bonding in ethene, we must first distinguish between two kinds of bonds.
A s (sigma) bond is a “head-to-head” overlap of orbitals with a cylindrical shape about the bond axis. This occurs when two “s” orbitals overlap or “p” orbitals overlap along their axis.
A p (pi) bond is a “side-to-side” overlap of parallel “p” orbitals, creating an electron distribution above and below the bond axis.
(See Animation: Carbon-Carbon Double Bond)
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69. (See Animation: Pi-Bond)
Figure 10.25
(See Animation: Pi-Bond) 2

70. Multiple Bonding
Now imagine that the atoms of ethene move into position.
Two of the sp2 hybrid orbitals of each carbon overlap with the 1s orbitals of the hydrogens.
The remaining sp2 hybrid orbital on each carbon overlap to form a s bond.
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71. Multiple Bonding
The remaining “unhybridized” 2p orbitals on each of the carbon atoms overlap side-to-side forming a p bond.
You therefore describe the carbon-carbon double bond as one s bond and one p bond.
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72. Molecular Orbital Theory
Molecular orbital theory is a theory of the electronic structure of molecules in terms of molecular orbitals, which may spread over several atoms or the entire molecule.
As atoms approach each other and their atomic orbitals overlap, molecular orbitals are formed.
In the quantum mechanical view, both a bonding and an antibonding molecular orbital are formed.
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73. Molecular Orbital Theory
For example, when two hydrogen atoms bond, a s1s (bonding) molecular orbital is formed as well as a s1s * (antibonding) molecular orbital. (animation only runs from CD)
The following slide illustrates the relative energies of the molecular orbitals compared to the original atomic orbitals.
Because the energy of the two electrons is lower than the energy of the individual atoms, the molecule is stable.
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74. H atom
H2 molecule
H atom
s1s* 1s 1s
s1s

75. Bond Order
The term bond order refers to the number of bonds that exist between two atoms.
The bond order of a diatomic molecule is defined as one-half the difference between the number of electrons in bonding orbitals, nb, and the number of electrons in antibonding orbitals, na.
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76. Electronic Configurations of Diatomic Molecules
In heteronuclear diatomic molecules, such as CO or NO, we must have additional molecular orbitals.
The overlap of “p” orbitals results in two sets of s orbitals (two bonding and two antibonding) and one set of p orbitals (one bonding and one antibonding). (See Animation: Pi Bond and Antibond).
The next slide illustrates the relative energies of these molecular orbitals.
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77. What Molecule does this diagram represent?
The arrows show the occupation of molecular orbitals by the valence electrons in N2. (See Animation: Molecular Orbital Diagram for a a Homonuclear Diatomic Molecule)
What Molecule does this diagram represent? 2

78. Practice Problems Chapter 10
27, 29, 37, 43, 53,
55, 56 (draw energy level diagrams, and calculate bond order)
67, 68, 71, 77, 73, 83, 88
I will do 79, as an example, class should draw MO diagrams for both products.
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79. Operational Skills Predicting molecular geometries.
Relating dipole moment and molecular geometry.
Applying valence bond theory.
Describing molecular orbital configurations.
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80. Animation: Electron Pair Repulsion, 2 Pairs
(Click here to open QuickTime animation)
Return to Slide 6
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81. Animation: Electron Pair Repulsion, 3 Pairs
(Click here to open QuickTime animation)
Return to Slide 6
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82. Animation: Electron Pair Repulsion, 4 Pairs
(Click here to open QuickTime animation)
Return to Slide 6
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83. Animation: Iodine Peutafluoride Structure
(Click here to open QuickTime animation)
Return to Slide 25
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84. Animation: Polar Molecules
(Click here to open QuickTime animation)
Return to Slide 28
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85. Return to Slide 27
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86. Figure 10.21: Bonding in HCl. Return to Slide 29
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87. Figure 10.22: Spatial arrangement of sp3 hybrid orbitals.
Return to Slide
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88. Figure 10.23: Diagrams of hybrid orbitals showing their spatial arrangements.
Return
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89. Animation: sp Hydridization
(Click here to open QuickTime animation)
Return to Slide 37
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90. Animation: sp2 Hydridization
(Click here to open QuickTime animation)
Return to Slide 37
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91. Animation: sp3 Hydridization
(Click here to open QuickTime animation)
Return to Slide 37
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92. Animation: Carbon-Carbon Double Bond
(Click here to open QuickTime animation)
Return to Slide 58
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93. (Click here to open QuickTime animation)
Animation: Pi-Bond
(Click here to open QuickTime animation)
Return to Slide 59
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94. Animation: sOrbitals/Bonding and Anti-Bonding
(Click here to open QuickTime animation)
Return to Slide 63
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95. Animation: Pi Bond and Antibond
(Click here to open QuickTime animation)
Return to Slide 66
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96. (Click here to open QuickTime animation)
Animation: Molecular Orbital Diagram for a Homonuclear Diatomic Molecule
(Click here to open QuickTime animation)
Return to Slide 67
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97. Fig. 10.8
Return
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98. Formation of a Sigma Bond
See H2 energy diagram animation
Also see s-orbital bonding with p-orbitals
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99. Return
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