1. Advanced Theories of Covalent Bonding
Valence Bond Theory
Hybridization of Atomic Orbitals
Multiple Bonds
Molecular Orbital Model

2. Valence Shell Electron Pair Repulsion (VSEPR) Theory
In molecules or polyatomic ions valence electrons exist in pairs;
Electron pairs repel each other:
they are arranged in an orientation that gives minimum electron-pair repulsions;
Relative strength of electron-pair repulsions:
Nonbonding pair repulsion > Nonbonding-bonding repulsion > bonding pair repulsion;

3. Determination of Molecular Shapes by the VSEPR Model
Molecular shapes determined by number and orientation of electron-pairs around central atom;
Shapes with minimum electron-pairs repulsion are preferred:
Nonbonding pairs placed farthest apart from each other.

4. Summary of Molecular Shapes predicted using the VSEPR Model
———————————————————————————————————————
# of E-pair # of # of Molecular
e-pairs Geometry bond-pairs Lone-pairs Shapes Examples
2 Linear Linear BeCl2
3 trigonal Trigonal BF3
planar planar
'' V-shape GeCl2
4 tetrahedral Tetrahedral CCl4
'' Trigonal
pyramidal PCl3
'' V-shape SCl2
5 trigonal Trigonal
bipyramidal bipyramidal AsCl5
'' See-saw SeCl4
'' T-shape ClF3
'' Linear KrF2
6 octahedral Octahedral SeCl6
'' Square-
pyramidal ICl5
'' Square
planar XeF4
——————————————————————————————————————

5. Molecular Shapes by VSEPR Model

6. How to Predict Polarity
Polarity of molecule depends on molecular shape and bond polarity.
Molecules (or polyatomic ions) with symmetrical shapes are non-polar;
Molecules with single lone-pairs on central atoms are generally polar;

7. Models for Chemical Bonding
Localized electron model
(Valence bond method and orbital hybridization)
Molecular orbital Theory
(Delocalized electron model)

8. Localized Electron Model
Good for explaining molecular shapes;
Two singly occupied orbitals from separate atoms form s-overlaps.
Electron pairs occupy overlap area to form s-bonds;
Two types of covalent bonds may be formed:
s -bond: orbitals overlap along internuclear axis
p -bond: side-to-side overlap of orbitals perpendicular to the internuclear axis

9. Orbital Hybridization
Hybridization describes the mixing of atomic orbitals on the central atom to account for the bond length, bond energy, and bond angle.

10. Consider bonding in CH4 C H
Facts: Methane (CH4) has tetrahedral shape with bond angles of 109.5o; all C-H bonds are identical.
If overlaps involve unhybridized orbitals, there would be one 2s-1s overlap and three 2p-1s overlaps.
One C-H bond in CH4 would be different from the others.
Since 2p orbitals are perpendicular to one another, bond angles in CH4 would be 90°.
Conclusion: overlap of orbitals in CH4 could not involve original atomic orbitals on carbon.

11. Overlaps between pure atomic orbitals lead to a molecular shape that does not agree with that predicted by VSEPR method or with observed shape. That is,
One C─H bond (from 2s-1s overlap) would be shorter;
Three C─H bonds (from 2p-1s overlaps) would be longer and at right angles to each other.
CH4 is tetrahedral and has symmetrical shape
Elements involved in bonding will modify their orbitals to achieve minimum energy configuration.

12. Hybridization of atomic orbitals in CH4
The 2s and 2p atomic orbitals in carbon mix to form 4 identical hybrid orbitals for bonding. Mixing of atomic orbitals is called hybridization.
Hybridization of one 2s and three 2p orbitals yields four sp3 hybrid orbitals on the carbon atom.
Four sp3 hybrid orbitals around a central atom yield a tetrahedral arrangement.
Each sp3 hybrid orbital overlap with 1s orbital of H-atom yields a tetrahedral molecule of CH4.

13. When a C-atom bonds to four other atoms, it hybridized its 2s and 2p atomic orbitals to form 4 new sp3 hybrid orbitals
The shape of each sp3 orbital shape is a mixture of one 2s and three 2p characters, yielding a large and a small lobes
Number of hybrid orbitals same as number of atomic orbitals hybridized.

14. All four sp3 orbitals have same energy level (degenerate)
H-atom can only use 1s orbital for bonding. The shared electron pair can be found in the overlap area of H1s—Csp3

15. NH3 and H2O also use sp3 hybridization; lone pairs occupy some of the hybrid orbitals.

16. Orbitals Hybridization
sp2 Hybridization
C2H4, ethylene is one example
Lewis and VSEPR structures tell us what to expect
H atoms still can only use 1s orbitals
C atom hybridizes 2s and two 2p orbitals into 3 sp2 hybrid orbitals

17. Formation of C=C Double Bond in Ethylene

18. The new sp2 orbitals are degenerate and in the same plane
One 2p orbital is unhybridized; it is perpendicular to the sp2 plane
One C—C bond in ethylene formed by overlap of sp2 orbitals, one from each C atoms, is called sigma (s) bond - the overlap is along the internuclear axis.
The second C—C bond formed by overlap of unhybridized 2p orbitals, is called pi (p) bond because the overlap is perpendicular to the internuclear axis.

19. 9. Pictorial views of the orbitals in ethylene

20. The Sigma System for Benzene

21. Resonance in Benzene

22. Delocalization Describes molecules that require resonance.
In molecules that require resonance, it is the  bonding that is most clearly delocalized, the  bonds are localized.
p orbitals perpendicular to the plane of the molecule are used to form  molecular orbitals.
The electrons in the  molecular orbitals are delocalized above and below the plane of the molecule.

23. The Pi System for Benzene

24. sp Hybridization
CO2, carbon dioxide is our example
Lewis and VSEPR predict linear structure
C atom uses 2s and one 2p orbital to make two sp hybrid orbitals that are 180 degrees apart
We get 2 degenerate sp orbitals and two unaltered 2p orbitals

25. Oxygen uses sp2 orbitals to overlap and form sigma bonds with C
Free p orbitals on the O and C atoms form pi bonds to complete bonding

26. d-orbitals can also be involved in hybridization
dsp3 hybridization in PCl5
d2sp3 hybridization in SF6

27. Types of Hybridization
There are five different types of hybridization:
sp - required by two pairs of electrons around an atom.
sp2 - required by three pairs of electrons around an atom.
sp3 - required by four pairs of electrons around an atom.
dsp3 - required by five pairs of electrons around an atom.
d2sp3 - required by six pairs of electrons around an atom.
A bond acts as one effective pair of electrons whether it is a single, double, or triple bond.

28. Hybridization and the Localized Electron Model
# Electron-pairs on Central Atom
Arrangement
Hybridization 2
linear sp 3
trigonal planar
sp2 4
tetrahedral
sp3 5
trigonal bipyramid
dsp3 or sp3d 6
octahedral
d2sp3 or sp3d2

29. F. The Localized Electron Model
1. Draw the Lewis structure(s)
2. Determine the arrangement of electron pairs (VSEPR model).
3. Specify the necessary hybrid orbitals.

30. Exercise Draw the Lewis structure for PCl5.
What is the shape of a phosphorus pentachloride molecule?
trigonal bipyramidal
What are the bond angles?
90o and 120o
The shape is trigonal bipyramidal with bond angles of 90o and 120o.

31. sp3d or dsp3 Hybridization
Hybridization of a 3s, three 3p, and a 3d orbitals yields five sp3d hybrid orbitals.
Five hybrid orbitals yield a trigonal bipyramidal arrangement.
If all five orbitals form covalent bonds, a molecule with trigonal bipyramidal shape is obtained.

32. The Orbitals used to form the Bonds in PCl5

33. Exercise Draw the Lewis structure for XeF4.
What is the shape of a xenon tetrafluoride molecule?
octahedral
What are the bond angles?
90o and 180o
The shape is octahedral with bond angles of 90o and 180o.

34. sp3d2 or d2sp3 Hybridization
Combination of one s, three p, and two d orbitals yields six sp3d2 hybrid orbitals;
Six equivalent sp3d2 hybrid orbitals around a central tom gives an octahedral arrangement.
If all six orbitals form covalent bonds, an octahedral molecule is obtained.

35. Hybridization of the Xenon Atom in XeF4

36. Concept Check-#1 Draw the Lewis structure for HCN.
Which hybrid orbitals are used?
Draw HCN:
Showing all bonds between atoms.
Labeling each bond as  or .
sp hybrid orbitals are used in the bonding of HCN. Two sigma bonds and two pi bonds are present.

37. Concept Check-#2 NH3 SO2 KrF2 CO2 ICl5 NH3 – 109.5o, sp3
Determine the bond angle and expected hybridization of the central atom for each of the following molecules:
NH3 SO2 KrF CO2 ICl5
NH3 – 109.5o, sp3
SO2 – 120o, sp2
KrF2 – 90o, 120o, dsp3
CO2 – 180o, sp
ICl5 – 90o, 180o, d2sp3
NH3 – 109.5o, sp3
SO2 – 120o, sp2
KrF2 – 90o, 120o, dsp3
CO2 – 180o, sp
ICl5 – 90o, 180o, d2sp3

38. Using the Localized Electron Model
Draw the Lewis structure(s).
Determine the arrangement of electron pairs using the VSEPR model.
Specify the hybrid orbitals needed to accommodate the electron pairs.

39. Exercise-#1: Hybridization
Describe the bonding in the water molecule using the localized electron model.
Solution
Since there are four electron pairs around oxygen atom, the oxygen is sp3 hybridized. The water molecule has a bent structure (or V-shape) O H

40. Sigma (s) and Pi (p) Bonds
Bonds formed by end-to-end overlaps of orbitals along the orbital axes are called sigma (σ) bonds.
Bonds formed by sidewise overlaps between two parallel unhybridized p-orbitals are called pi (π) bonds. The overlaps of orbitals occur above and below the bond axis
All single bonds are σ-bonds.
A double bond consists of one σ bond and one π bond.
A triple bond is made up of one σ and two π bonds.

41. Exercise-#2: Sigma and Pi Bonds
How many σ bonds are there in the commercial insecticide, “Sevin,” shown below? How many π bonds?
Solution
There are 27 σ bonds.
There are 6 π bonds.

42. Practice with Hybrid Orbitals
Give the hybridization and predict the geometry of the central atom in the ion IF2+.
Lewis structure:
There are 4 effective electron pairs around the central atom making it sp3 hybridized.
The VSEPR structure has a tetrahedral basis but since it has two bonding pairs, it will take on a bent shape with a bond angle smaller than 109.5°. F I

43. Molecular Shape of SOF4 SOF4
Sulfur in SOF4 has 5 electron pairs. The VSEPR structure is a trigonal bipyramid. Sulfur is sp3d hybridized. S F O

44. Hybridization and Structure of C2H2
HC≡CH
Each carbon has 2 effective electron pairs. The VSEPR structure is linear. Each carbon is sp hybridized.
C C H

45. The Molecular Orbital Theory
The localized electron model, which includes orbital hybridization, does an excellent job in predicting and justifying molecular shapes, but it does not explain why certain molecules are paramagnetic.
The molecular orbital (MO) theory views electrons as belonging to the entire molecule rather than to individual atoms in the molecule.
MO theory provides explanation why molecules, such as B2 and O2 are paramagnetic, where as C2 and N2 are diamagnetic.

46. Molecular Orbital Theory
Regards a molecule as a collection of nuclei and electrons, where the electrons are assumed to occupy orbitals that extend over the entire molecule.
Electrons occupy molecular orbitals in the same manner they do the atomic orbitals.
The electrons are assumed to be delocalized over the entire molecule, rather than always located between a pair of atoms.

47. Electron Probability in MO
The electron probability of both molecular orbitals is centered along the line passing through the two nuclei.
Sigma (σ) molecular orbitals (MOs)
In the molecule only the molecular orbitals are available for occupation by electrons.

48. Key Ideas of the MO Model
All valence electrons in a molecule exist in a set of molecular orbitals.
Valence orbitals of each atom are not acting independently, but rather interact as a whole to form a set of molecular orbitals.
Each pair of atomic orbitals interacts to form a set of bonding and antibonding molecular orbitals.
Energy for bonding molecular orbitals is lower than their corresponding atomic orbitals.
While antibonding molecular orbitals have higher energy than their atomic orbitals.

49. Combination of Hydrogen 1s Atomic Orbitals to form MOs

50. Bonding in H2

51. Sigma Bonding and Antibonding Molecular Orbitals

52. (a) MO Energy-Level Diagram for H2 Molecule (b) Shapes of MOs

53. Molecular Orbital Energy Levels
The sns -bonding molecular orbital (MO1) - lower energy than the ns atomic orbital;
The sns*-antibonding molecular orbital (MO2) - higher energy than the ns atomic orbital

54. MO Energy-Level Diagram for the H2 Molecule

55. Example: H2

56. Bond Order in MO Model
BO =
# bonding electrons - # antibonding electrons 2
Bond order (BO) is a measure of net bonding interactions.
BO must be greater than 0 for a stable molecule to form.
The higher the BO, the stronger the bond.

57. Bond Order in H2–

58. Bond Order
Larger bond order means greater bond strength.

59. MO Energy-Level Diagram for He2 Bond order = 0 (molecule cannot form)

60. Three Mutually Perpendicular 2p Orbitals on Two Adjacent Atoms of B - Ne

61. Constructive Interference from Combining Parallel p Orbitals

62. Set of π MOs

63. Pi Bonding and Antibonding Orbitals

64. Expected MO Energy-Level Diagram from Combination of 2p Orbitals on Boron

65. Expected MO Energy-Level Diagram for B2 Molecule

66. The Corrected MO Energy-Level Diagram for B2 Molecule

67. MO Diagrams for Homonuclear Diatomic Molecules in the Second Period.

68. MO Energy-level Diagram for B2
Filling molecular orbitals s2s and s2p works the same as when filling s1s orbitals.
Fill degenerate orbitals p2p separately, then pair-up with electrons having opposite spins.
Molecular orbital diagram shows that B2 molecule is paramagnetic.
BO in B2 = ½ (4 – 2) = 1
σ*2s
σ2s E
π2p
σ2p
π*2p
σ*2p

69. MO Energy-level Diagram for C2
Molecular orbital diagram for C2 shows that the molecule is diamagnetic.
Bond order in C2
= ½ (6 – 2) = 2
σ*2s
σ2s E
π2p
σ2p
π*2p
σ*2p

70. MO energy-level diagram for N2
Molecular orbital diagram for N2 shows that it is diamagnetic.
Bond order in N2
= ½ (8 – 2) = 3
σ*2s
σ2s E
π2p
σ2p
π*2p
σ*2p

71. MO Energy-level Diagram for O2
It shows that O2 molecules is paramagnetic
Bond order in O2 :
½ (8 – 4) = 2
σ*2s
σ2s E
π2p
σ2p
π*2p
σ*2p

72. Molecular Orbital Summary of Second Row Diatomic Molecules

73. Electron Configurations in MO
Electron configuration of a diatomic molecule is written in a similar manner as that of the atom:
Each molecular orbital can hold 2 electrons with opposite spins.
Examples:
N2: (s1s)2 (s*1s)2 (s2s)2 (s*2s)2 (p2s)4 (s2p)2
O2: (s1s)2 (s*1s)2 (s2s)2 (s*2s)2 (s2p)2 (p2s)4 (s*2p)2

74. Liquid Oxygen Poured into Space Between Poles of Magnet
Donald W. Clegg

75. Paramagnetism
Paramagnetism – substance is attracted into the inducing magnetic field.
Unpaired electrons (O2)
Diamagnetism – substance is repelled from the inducing magnetic field.
Paired electrons (N2)

76. Apparatus used to Measure the Paramagnetic Property of Compounds

77. Paramagnetism vs. Diamagnetism
Paramagnetism implies the presence of unpaired electrons in an atom or molecule; diamagnetism implies all electrons in an atom or molecule are paired.
Paramagnetic substances are attracted to a magnet; whereas diamagnetic substances are repelled by magnet.
Paramagnetic effect is generally much stronger than diamagnetic effect.

78. MO Energy-Level Diagram for NO Molecule

79. MO Energy-Level Diagram for NO+ and CN- Ions

80. Homonuclear Diatomic Molecules
Composed of 2 identical atoms.
Only the valence orbitals of each atom contribute significantly to the molecular orbitals of a particular molecule.

81. Heteronuclear Molecules
Molecules made up of different atoms.
Molecules that are made up of atoms close together in the periodic table have similar molecular orbital diagrams as homonuclear molecules.
When two atoms of a diatomic molecule are very different, the energy-level diagram is more complicated.

82. Heteronuclear Diatomic Molecule: HF
The 2p atomic orbitals in fluorine has lower energy than the 1s atomic orbital in hydrogen;
Fluorine nucleus has 9 protons, which binds valence electrons more tightly than hydrogen nucleus.
In HF molecule, electron pair in s-bonding molecular orbital is closer to the fluorine atom, resulting in a polar molecule;

83. The Electron Probability Distribution in the Bonding Molecular Orbital of the HF Molecule

84. Orbital Energy-Level Diagram for the HF Molecule

85. Molecular Orbital Energy Diagram for HF
Because the fluorine 2p orbital is lower in energy than the hydrogen 1s orbital, electrons tend to be closer to the fluorine atom.

86. Exercise #2 on Molecular Orbitals
Determine the electron configuration and bond order for the following. If it exists, discuss its magnetism.
1. F2-
σ*2s
σ2s E
π2p
σ2p
π*2p
σ*2p
How many valence electrons does F2- have?
15 valence electrons
BO =
8 b.e. - 7 a.e. 2
= ½
F2- would be paramagnetic.

87. Exercise #3 on Molecular Orbitals
Determine the electron configuration and bond order for the following. If it exists, discuss its magnetism.
2. S22-
σ*3s
σ3s E
π3p
σ3p
π*3p
σ*3p
How many valence electrons does S22- have?
14 valence electrons
BO =
8 b.e. - 6 a.e. 2
= 1
S22- would be diamagnetic.

88. Exercise #4 on Molecular Orbitals
Determine the electron configuration and bond order for the following. If it exists, discuss its magnetism.
3. Cl22+
σ*3s
σ3s E
π3p
σ3p
π*3p
σ*3p
How many valence electrons does Cl22+ have?
12 valence electrons
BO =
8 b.e. - 4 a.e. 2
= 2
Cl22+ would be paramagnetic.

89. Molecular Orbital versus Localized Electron Model
Advantage of molecular orbital model: it correctly predicts relative bond strength and magnetism of simple diatomic molecules.
It explain bond polarity base on orbital energy diagram and electron probability; portrays electrons as being delocalized in polyatomic molecules.
Major disadvantage: energy diagram too complex for polyatomic molecules; an approximation is used.

90. Molecular Orbital versus Localized Electron Model
The localized electron model assumes that electrons are confined between a given pair of atoms in a molecule.
Combining the localized electron and molecular orbital models gives us a more accurate description, such that:
Electrons in σ-bonds are considered localized.
Electrons involved in p-bonds are said to be delocalized.

91. Delocalized Pi Electrons in Benzene