1. Chapter 9 Molecular Geometries and Bonding Theories
Chemistry, The Central Science, 10th edition
Theodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten
Chapter 9 Molecular Geometries and Bonding Theories
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice-Hall, Inc.

2. What Determines the Shape of a Molecule?
The Lewis structure is drawn with the atoms all in the same plane.
The overall shape of the molecule will be determined by its bond angles.

3. What Determines the Shape of a Molecule?
Electron pairs, whether they be bonding or nonbonding, repel each other.
By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

4. For molecules whose formula is of the general form ABn, there are five fundamental shapes.

5. Electron Domains
We can refer to the electron pairs as electron domains.
Each pair of electrons count as an electron domain, whether they are in a lone pair, in a single, double or triple bond.
This molecule has four electron domains around the central atom.

6. Electron-Domain Geometries
Count the number of electron domains around the central atom in the Lewis structure.
The geometry will be that which corresponds to that number of electron domains.

7. In order to predict molecular shape, we assume the valence electrons repel each other. Therefore, the molecule adopts whichever 3D geometry minimizes this repulsion.
We call this process Valence Shell Electron Pair Repulsion (VSEPR) theory.
There are simple shapes for AB2 to AB6 molecules.

8. Electron domain geometry: When considering the geometry about the central atom, we consider all electrons (lone pairs and bonding pairs).
When naming the molecular geometry, we focus only on the positions of the atoms.

9. VSEPR Model
To determine the shape of a molecule, we distinguish between lone pairs (or non-bonding pairs, those not in a bond) of electrons and bonding pairs (those found between two atoms).
We define the electron domain geometry (or orbital geometry) by the positions in 3D space of ALL electron pairs (bonding or non-bonding).
The electrons adopt an arrangement in space to minimize e--e- repulsion.

12. Examples – Draw the Lewis structures, and then determine the orbital geometry of each. Indicate the number of electron domains first:
H2S
CO2
PCl3
CH4
SO2

13. Examples – Draw the Lewis structures, and then determine the orbital geometry of each:
e- domain orbital geometry
or e- domain geom
H2S tetrahedral
CO linear
PCl tetrahedral
CH tetrahedral
SO trigonal planar

14. To determine the electron pair ( electron domain) geometry:
draw the Lewis structure,
count the total number of electron pairs around the central atom,
arrange the electron pairs in one of the above geometries to minimize e--e- repulsion, and count multiple bonds as one bonding pair.
But then we have to account for the shape of the molecule

15. Molecular Geometries
Within each electron domain, then, there might be more than one molecular geometry.

16. Linear Electron Domain
In this domain, there is only one molecular geometry: linear.
NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

17. Nonbonding Pairs and Bond Angle
Nonbonding pairs are physically larger than bonding pairs.
Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule.

18. By experiment, the H-X-H bond angle decreases on moving from C to N to O:
Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs.
Therefore, the bond angle decreases as the number of lone pairs increase.

19. Similarly, electrons in multiple bonds repel more than electrons in single bonds.

20. Multiple Bonds and Bond Angles
Double and triple bonds place greater electron density on one side of the central atom than do single bonds.
Therefore, they also affect bond angles.

21. Trigonal Planar Electron Domain
There are two molecular geometries:
Trigonal planar, if all the electron domains are bonding
Bent, if one of the domains is a nonbonding pair.

22. Tetrahedral Electron Domain
There are three molecular geometries:
Tetrahedral, if all are bonding pairs
Trigonal pyramidal if one is a nonbonding pair
Bent if there are two nonbonding pairs

24. Examples – Same examples as before, now determine the the molecular geometry of each, including shapes and bond angles:
H2S
CO2
PCl3
CH4
SO2

25. Examples – Determine the molecular geometry of each, including shapes and bond angles:
Shape Angles
H2S bent <109°
CO2 linear 180°
PCl3 trigonal pyramid <109°
CH4 tetrahedral °
SO2 bent <120°

26. Molecules with Expanded Valence Shells
For elements of the 3rd shell and below, some atoms can have expanded octets.
AB5 (trigonal bipyramidal) or AB6 (octahedral) electron pair geometries.

27. Trigonal Bipyramidal Electron Domain (5 e domains)
There are two distinct positions in this geometry:
Axial
Equatorial

28. Molecules with Expanded Valence Shells
To minimize e--e- repulsion, lone pairs are always placed in equatorial positions.

29. Trigonal Bipyramidal (e- domain)
There are four distinct molecular geometries in this domain:
*Trigonal bipyramidal
*Seesaw
*T-shaped
*Linear

31. Trigonal Bipyramidal Electron Domain
Lower-energy conformations result from having nonbonding electron pairs in equatorial, rather than axial, positions in this geometry.

32. Octahedral electron domain
6 electron pairs
All positions are equivalent in the octahedral domain.
There are three molecular geometries:
*Octahedral
*Square pyramidal
*Square planar

35. Examples – Determine the Shape of each, indicate the electron domain, molecular geometry and angles.
PF5
XeF4
SF6
SCl4

36. Examples – Determine the Shape of each:
PF5 trigonal bipyramid 90°, 120°
XeF4 square planar °
SF6 octahedral °
SCl4 see-saw <90°,< 120°
See moving chart

37. Larger Molecules
In larger molecules, it makes more sense to talk about the geometry about a particular atom rather than the geometry of the molecule as a whole.

38. Larger Molecules
This approach makes sense, especially because larger molecules tend to react at a particular site in the molecule.

39. Shapes of Larger Molecules
In acetic acid, CH3COOH, there are three central atoms.
We assign the geometry about each central atom separately.

40. Examples – Determine the shape and angles about each atom:1 2

41. Examples – Determine the shape and angles about each atom:1 2
Trigonal planar linear
Bent trigonal pyramid
Trigonal planar

42. Molecular Shape and Molecular Polarity
When there is a difference in electronegativity between two atoms, then the bond between them is polar.
It is possible for a molecule to contain polar bonds, but not be polar.
For example, the bond dipoles in CO2 cancel each other because CO2 is linear.

44. In water, the molecule is not linear and the bond dipoles do not cancel each other.
Therefore, water is a polar molecule.
The overall polarity of a molecule depends on its molecular geometry.

46. Polarity
By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule.

47. Polarity

48. ABn Nonpolarity
For ABn molecules in which all the B atoms are the same, certain symmetrical shapes lead to nonpolar molecules
Linear (AB2), trigonal planar (AB3), tetrahedral and square planar (AB4), trigonal bipyramidal (AB5) and octahedral (AB6) are all nonpolar

49. Examples – Determine whether each is polar or nonpolar:
CCl4
PCl3
BF3
BrClFCH
SO3

50. Examples – Determine whether each is polar or nonpolar:
CCl4 nonpolar
PCl3 polar
BF3 nonpolar
BrClFCH polar
SO3 nonpolar

51. Overlap and Bonding
We think of covalent bonds forming through the sharing of electrons by adjacent atoms.
In such an approach this can only occur when orbitals on the two atoms overlap.

52. Overlap and Bonding
Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electron-electron repulsion.
However, if atoms get too close, the internuclear repulsion greatly raises the energy.

53. HYBRIDIZATION
VALENCE BOND THEORY (s and p bonds)
BOND ORDER

54. Hybrid Orbitals
It’s hard to imagine tetrahedral, trigonal bipyramidal, and other geometries arising from the atomic orbitals we recognize.

55. Hybrid Orbitals
Atomic orbitals can mix or hybridize in order to adopt an appropriate geometry for bonding.
Hybridization is determined by the electron domain geometry.
sp Hybrid Orbitals
Consider the BeF2 molecule (experimentally known to exist):

56. Hybrid Orbitals Consider beryllium:
In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals.

57. Hybrid Orbitals
But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds.

58. Hybrid Orbitals
Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals.
These sp hybrid orbitals have two lobes like a p orbital.
One of the lobes is larger and more rounded as is the s orbital.

59. Hybrid Orbitals (sp)
These two degenerate orbitals would align themselves 180 from each other.
This is consistent with the observed geometry of beryllium compounds: linear.

60. Hybrid Orbitals
With hybrid orbitals the orbital diagram for beryllium would look like this.
The sp orbitals are higher in energy than the 1s orbital but lower than the 2p.

62. Since only one of the Be 2p orbitals has been used in hybridization, there are two unhybridized p orbitals remaining on Be.

63. sp2 Hybrid Orbitals
Important: when we mix n atomic orbitals we must get n hybrid orbitals.
sp2 hybrid orbitals are formed with one s and two p orbitals. (Therefore, there is one unhybridized p orbital remaining.)
The large lobes of sp2 hybrids lie in a trigonal plane.
All molecules with trigonal planar electron pair geometries have sp2 orbitals on the central atom.

64. Hybrid Orbitals (sp2)
For boron

65. Hybrid Orbitals
…three degenerate sp2 orbitals.

66. November 29 HYBRIDIZATION WITH MULTIPLE BONDS – VALENCE BOND THEORY
ORDER OF A BOND
PARAMAGNETIC DIAMAGNETIC MOLECULES
PRACTICE PROBLEMS

67. Hybrid Orbitals
With carbon we get…

68. Hybrid Orbitals
…four degenerate
sp3 orbitals.

69. sp2 and sp3 Hybrid Orbitals
sp3 Hybrid orbitals are formed from one s and three p orbitals. Therefore, there are four large lobes.
Each lobe points towards the vertex of a tetrahedron.
The angle between the large lobs is 109.5.
All molecules with tetrahedral electron pair geometries are sp3 hybridized.

70. Hybridization Involving d Orbitals
Since there are only three p-orbitals, trigonal bipyramidal and octahedral electron domain geometries must involve d-orbitals.
Trigonal bipyramidal electron domain geometries require sp3d hybridization.
Octahedral electron domain geometries require sp3d2 hybridization.
Note the electron domain geometry from VSEPR theory determines the hybridization.

71. Sp3d and sp3d2 Hybrid Orbitals
For geometries involving expanded octets on the central atom, we must use d orbitals in our hybrids.

72. Hybrid Orbitals This leads to five degenerate sp3d orbitals…
…or six degenerate sp3d2 orbitals.

73. Hybrid Orbitals Summary Draw the Lewis structure.
Determine the electron domain geometry with VSEPR.
Specify the hybrid orbitals required for the electron pairs based on the electron domain geometry.

76. Hybrid Orbitals
Once you know the electron-domain geometry, you know the hybridization state of the atom.

77. Examples – Determine the hybridization on the central atom of each:
NCl3
CO2
H2O
SF4
BF3
XeF4

78. Examples – Determine the hybridization on the central atom of each:
NCl3 sp3
CO2 sp
H2O sp3
SF4 sp3d
BF3 sp2
XeF4 sp3d2

79. Examples – Determine the hybridization on EACH atom:

80. Examples – Determine the hybridization on EACH atom:
sp2 sp
sp3
sp3
sp2

81. Valence Bond Theory
Hybridization is a major player in this approach to bonding.
There are two ways orbitals can overlap to form bonds between atoms.

82. Sigma () Bonds Sigma bonds are characterized by Head-to-head overlap.
Cylindrical symmetry of electron density about the internuclear axis.

83. Pi () Bonds Pi bonds are characterized by Side-to-side overlap.
Electron density above and below the internuclear axis.

84. Single Bonds
Single bonds are always  bonds, because  overlap is greater, resulting in a stronger bond and more energy lowering.

85. Multiple Bonds
In a multiple bond one of the bonds is a  bond and the rest are  bonds.

86. Multiple Bonds
In a molecule like formaldehyde (shown at left) an sp2 orbital on carbon overlaps in  fashion with the corresponding orbital on the oxygen.
The unhybridized p orbitals overlap in  fashion.

87. Multiple Bonds
In triple bonds, as in acetylene, two sp orbitals form a  bond between the carbons, and two pairs of p orbitals overlap in  fashion to form the two  bonds.

88. Delocalized Electrons: Resonance
When writing Lewis structures for species like the nitrate ion, we draw resonance structures to more accurately reflect the structure of the molecule or ion.

89. Delocalized Electrons: Resonance
In reality, each of the four atoms in the nitrate ion has a p orbital.
The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen.

90. Delocalized Electrons: Resonance
This means the  electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion.

91. Resonance
The organic molecule benzene has six  bonds and a p orbital on each carbon atom.

92. Resonance
In reality the  electrons in benzene are not localized, but delocalized.
The even distribution of the  electrons in benzene makes the molecule unusually stable.

93. General Conclusions
Every two atoms share at least 2 electrons.
Two electrons between atoms on the same axis as the nuclei are  bonds.
-Bonds are always localized.
If two atoms share more than one pair of electrons, the second and third pair form -bonds.
When resonance structures are possible, delocalization is also possible.

94. Bond Order = total number of covalent bonds between 2 atoms
Bond order = 1 for single bond.
Bond order = 2 for double bond.
Bond order = 3 for triple bond.
If there is resonance then divide the bond order by the number of atoms that share the resonance structure.
Fractional bond orders are possible (with resonance)

95. Examples – Draw Lewis Structures (including resonance)
Examples – Draw Lewis Structures (including resonance). Determine the total number of σ and π bonds in the molecule, and the bond order of each bond: O3
SO3
CO2

96. Examples – Draw Lewis Structures (including resonance)
Examples – Draw Lewis Structures (including resonance). Determine the total number of σ and π bonds in the molecule, and the bond order of each bond: O3
SO3
CO2
2σ and 1π
Each bond 1.5
3σ and 1π each bond 1.33
2σ and 2π each bond 2

97. Electron Configurations and Molecular Properties
Two types of magnetic behavior:
paramagnetism (unpaired electrons in atom or molecule): strong attraction between magnetic field and molecule;
diamagnetism (no unpaired electrons in atom or molecule): weak repulsion between magnetic field and molecule.
Magnetic behavior is detected by determining the mass of a sample in the presence and absence of magnetic field:

98. Second-Row Diatomic Molecules
Electron Configurations and Molecular Properties
large increase in mass indicates paramagnetism,
small decrease in mass indicates diamagnetism.