1. Molecular Shapes, Valence Bond Theory, and Molecular
Lecture Presentation
Chapter 10 Chemical Bonding II:
Molecular Shapes, Valence Bond Theory, and Molecular
Orbital Theory
Catherine MacGowan
Armstrong Atlantic State University
© 2013 Pearson Education, Inc.

2. MOLECULAR GEOMETRY AND VSEPR
Valence Shell Electron
Pair Repulsion
VSEPR is used to predict the molecular geometries of compounds based on the number of electron groups attached to the center atom of the molecule.
Most important factor in determining the geometry is minimizing the relative repulsion between electron pairs.
Molecules’ or ions’ geometry reflects arrangement that favors minimal electron repulsion.

3. Electron Pair Geometries: Family Groups

4. Structure Determines Properties
The chemical and physical properties of molecular substances depend on the structure of the species.
Factors to be considered:
Skeletal arrangement of the atoms
Who is bonded to who
Type of intramolecular bonding between the atoms
ionic, polar covalent, or covalent
Shape of the molecule
Bonding theory should allow you to predict the shapes of molecules.

5. Molecular Geometry
Molecules are three-dimensional objects, so their shape can be described in terms related to geometric figures.
Example:
Linear
Tetrahedral
Octahedral
Geometric shapes have characteristic:
“corners” that indicate the positions of the surrounding atoms around a central atom in the center of the geometric figure, and
angles (e.g., bond angles)

6. Lewis Structures and VSEPR Theory
A Lewis structure
Predicts the number of valence electron pairs around the central atom(s)
Each lone pair of electrons constitutes one electron group on a central atom.
Each bond constitutes one electron group on a central atom.
regardless of whether it is a single, double, or triple bond

7. Lewis Structures and VSEPR Theory
Based on the number of electron groups attached to the center atom of the molecule
The arrangement of the atoms attached to the center atom minimizing the relative repulsion between electron pairs
Note: It is the arrangement of atoms attached to the center atom that dictates the compound’s structure (electron pair geometry) and shape (molecular geometry of atoms).

8. Geometric Arrangement of Electron Groups Around Center Atom: Electron Group Geometries
Five types of arrangements (families) of electron groups around a central atom
Types of electron groups
Atoms
Lone pairs
The five geometric arrangements (families)
Linear
Trigonal planar
Tetrahedral
Trigonal bipyramidal
Octahedral

9. Geometric Arrangement of Electron Groups Around Center Atom: Electron Group Geometries
Linear
Two electron groups attached to the center atom of a two-atom molecule
Bond angle between electron groups and center atom: 180o

10. Geometric Arrangement of Electron Groups Around Center Atom: Electron Group Geometries
Trigonal planar
Three electron groups attached to the center atom with each electron group occupying a corner position on a triangle
Bond angle between electron groups: 120o

11. Geometric Arrangement of Electron Groups Around Center Atom: Electron Group Geometries
Tetrahedral
Four electron groups surrounding the center atom with electron groups occupying corner positions of a tetrahedron
Bond angle between electron groups: 109.5o

12. Geometric Arrangement of Electron Groups Around Center Atom: Electron Group Geometries
Trigonal bipyramidal
Five electron groups attached to the center atom
The electron groups will occupy positions at the axial position (above and below the center atom) and at the equatorial positions in the base plane of the center atom.
Bond angle between equatorial positions is 120°.
Bond angle between axial and equatorial positions is 90°.

13. Geometric Arrangement of Electron Groups Around Center Atom: Electron Group Geometries
Octahedral
Six electron groups attached or surrounding the center atom
Electron groups in the shape of two square base pyramids that are base to base, with the central atom in the center of the shared bases
It is called octahedral because the geometric figure has eight sides.
All positions are equivalent.
The bond angle is 90°.

14. Geometric Derivations or Why Distortions Occur
Why do distortions in a molecule’s three-dimensional structure occur?
Lone pairs
Require more space
Need for space causes a distortion in the bond angle between the atoms.
Types of atoms bonding to the center atom
Size of atom
Types of bonds between the bonding groups.
Single vs. multiple
What results?
The molecular geometry (shape) will vary
Shape will be a derivative of one of the basic shapes.

15. Chart of Electron and Molecular Geometries

16. Chart of Electron and Molecular Geometries

17. Geometric Derivations and Placement of Lone Pairs
Placement of lone pairs around a center atom:
Look at the geometry of the molecule.
Want to minimize the electron-electron repulsion
Want to occupy a position within the structure that has space to allow for “spreading out”
Geometries that DO NOT have ‘”minimizing” positions around the center atom or where all bond angles are equally spaced:
Linear, trigonal planar, tetrahedral, and octahedral

18. The Effect of Lone Pairs on a Compound’s Molecular Geometry (Shape)
Lone pair/nonbonding electron groups “occupy more space” on the central atom.
Why?
Electron density is exclusively on the central atom rather than shared like bonding electron groups.
This affects the bond angles, making them smaller than expected.
The relative sizes of repulsive force interactions are:
(Lone Pair – Lone Pair) > (Lone Pair – Bonding Pair )> (Bonding Pair – Bonding Pair )

19. Distortion of Trigonal Planar Geometric Group
Happens when one of the electron groups is a lone pair
One lone pair + 2 atoms attached to center atom
Electron geometry is trigonal planar but the shape of molecule (molecular geometry) is BENT.
The bond angle between groups is <120°.
:O = S - O: :O - S = O: :O = S = O:
. .
. .
. .
. .
. .
. .
. .
. .
. .
. .
. .

20. H2O: Tetrahedral Bent Shape
Bonding pair
H : O : H
Lone pair
. .
. .

21. Distortion of Tetrahedral Geometric Group
It happens when ONE or TWO lone pairs surround the center atom.
All belong to the tetrahedral electron geometric group but will have different molecular geometries (shape).
Pyramidal:
1 lone pair surrounding center atom + 3 atom groups attached to center atom
Bent or V shape:
2 lone pairs surrounding center atom + 2 atom groups attached to the center atom
When there are 4 electron groups around the central atom
Bond angle for both distorted tetrahedral shapes is <109.5°.

22. Trigonal Bipyramidal Geometry and the Placement of Lone Pairs
Positions of trigonal bipyramidal structure:
Equatorial
BOND angle between equatorial positions is <120°.
Axial
bond angle between axial and equatorial positions is <90°.
Lone pairs will occupy the equatorial positions because there is more room. .

23. Distortion of Trigonal Bipyramidal Geometric Group
It happens when ONE, TWO, or THREE lone pairs surround the center atom.
All belong to the trigonal bipyramidal electron geometric group but will have different molecular geometries (shapes).
See Saw or Sawhorse:
1 lone pair at one of equatorial positions surrounding center atom + 4 atom groups attached to center atom

24. Distortion of Trigonal Bipyramidal Geometric Group
It happens when ONE, TWO, or THREE lone pairs surround the center atom.
All belong to the trigonal bipyramidal electron geometric group but will have different molecular geometries (shapes).
T shape:
2 lone pairs at the equatorial positions surrounding center atom + 3 atom groups attached to the center atom

25. Distortion of Trigonal Bipyramidal Geometric Group
It happens when ONE, TWO, or THREE lone pairs surround the center atom.
All belong to the trigonal bipyramidal electron geometric group but will have different molecular geometries (shapes).
Linear shape:
3 lone pairs at the equatorial positions surrounding center atom + 2 atom groups attached to the center atom

26. Distortion of the Octahedral Geometric Group
Six electron groups (atoms and/or lone pairs) surrounding the central atom indicates an octahedral electron geometry (family).
Distortions happen when one or more lone pairs occupy the six positions.
All positions in this geometry have the same bond angle (90o) when atom groups occupy the sites.
Therefore, no site is more favorable when substituting lone pair groups for atoms.
In distorted geometric structures the bond angle between axial and equatorial positions will be <90°.

27. Distortion of the Octahedral Geometric Group
It happens when ONE or TWO lone pairs surround the center atom.
All belong to the octahedral electron geometric group but will have different molecular geometries (shapes).
Square pyramidal:
1 lone pair at either the equatorial or axial position surrounding center atom + 5 atom groups attached to the center atom

28. Distortion of the Octahedral Geometric Group
It happens when ONE or TWO lone pairs surround the center atom.
All belong to the octahedral electron geometric group but will have different molecular geometries (shapes).
Square planar:
2 lone pairs at either the equatorial or axial positions surrounding center atom + 4 atom groups attached to the center atom

29. Geometric Arrangements and Resonance Structures
To achieve the best structure (energy-wise), a compound’s electron groups must be bonds and all the bonds must be equivalent.
For molecules that exhibit resonance, the molecular geometry will be the same for all possible arrangements.
Example: NO3-
Electron geometry (family) is trigonal planar.
Molecular geometry (shape) is also trigonal planar.

30. Predicting the Electron (Family Group) and Molecular (Shape) Geometries of Compounds
Draw the Lewis structure of the compound.
Determine the number of bonding groups (atoms bonded to center atom) and lone pairs surrounding the central atom.
Bonding groups (atoms bonded or attached to center atom):
You are counting the number of atoms bonded to the center atom and NOT the number of bonds between the atoms.
An atom double bonded to the center atom counts as ONE attachment.
Determine the geometries.
Electron (family group):
TOTAL the number of atoms (attachments) to the center atom + the number of lone pairs surrounding the center atom.
Example: If there are 3 atoms bonded (attached) to the center atom plus 2 lone pairs surrounding the center atom, the TOTAL number is 5.
5 = trigonal bipyramidal geometry
Molecular (shape):
The electron geometry dictates the molecular geometry associated with that geometric grouping.
Example: If the compound belongs to the trigonal bipyramidal (total of 5) and has 3 bonded atoms (attachments) and 2 lone pairs surrounding the center atom, then the molecular shape of the molecule is T-shape.

31. Problem: Determine the electron and molecular geometry for the compound ammonia, NH3.
Draw electron dot structure.
Count # of attachments and lone pairs on center atom.
Three attachments (hydrogen atoms) to the center atom (nitrogen) and one lone pair = 4
3. Electron pair geometry (family) for combination of 4 is tetrahedral .
4. The molecular geometry (shape) is trigonal pyramidal.

32. Problem: Determine the electron geometry (family) and molecular geometry (shape) for the molecule ClO2F.

33. Problem: Determine the electron geometry (family) and molecular geometry (shape) for the molecule BF3.
Answer:
B is the least electronegative of atoms in the molecule so it will be the center atom.
All atoms will be attached to the B atom.
Draw the Lewis structure:
B = e─
F = 3 x 7 e e─
Total = e─
3 electron groups on the center atom B
Electron geometry group (family)
is trigonal planar.
All groups surrounding the center atom
are atom groups (no lone pairs);
therefore ,the molecular geometry (shape)
of BF3 is trigonal planar.

34. Not-Quite-Perfect Geometry
NOTE:
Because the bonds are not identical, the observed angles are slightly different from the ideal.

35. Multiple Central Atoms and Their Geometries
Large structured molecules have more than ONE center atom
implying that they have multiple central atoms
Molecules with more than one center atom, therefore, have more than one type of geometry.
Geometries (electron and molecular) are assigned for each center atom in the molecule.

36. VESPR and Multi-Center Molecules
Methanol, CH3OH
Methanol has two center atoms.
Center one:
H-C-H = 109o
Electron geometry: Tetrahedron
Molecular geometry: Tetrahedral
Center two:
- C-O-H <109o
Molecular geometry: Bent

37. Representing Three-Dimensional Shapes on a Two-Dimensional Surface
It is difficult to draw molecules showing their dimensionality.
Convention dictates that the central atom is put in the plane of the paper.
To indicate on paper the spatial relationship of the atom groups and/or lone pairs to the center atom, the following guidelines and symbols are used:
Put as many other atoms as possible in the same plane and indicate with a straight line.
For atoms in front of the plane, use a solid wedge.
For atoms behind the plane, use a hashed wedge.

39. Molecular Polarity
Question: Why are some molecules polar in nature and others nonpolar?
Answer: Look at the bonding and their three-dimensional structure (geometry).
Molecules to be polar need to have:
Polar covalent bonds within the molecule
Difference in bond atoms’ electronegativity
A net dipole moment because the molecule’s electron geometry is NOT “symmetric”
Symmetric molecules DO NOT have a dipole moment.
The molecule is NONPOLAR.
Unsymmetric molecules DO have a dipole moment.
The molecule is POLAR.
Polarity affects the intermolecular forces of attraction.
Boiling points and solubilities
Like dissolves like
Nonbonding pairs affect molecular polarity; they pull the electron density strongly.

40. Dipole Moment and Polarity

41. Polar or Nonpolar?
Which one is polar: CO2 or H2O?
H2O CO2

42. Molecular Polarity Explains solubility of substances
Likes dissolve in likes.
Provides information on a substance’s chemical and physical properties
Melting points and boiling points
H2O vs CH4
Bp: 100 oC Bp: -161 oC

43. Molecular shape in food science:the production of artificial sweeteners(人工甘味劑), e.g., aspartame (200 times more sweet than sugar) and saccharin.
Artificial sweeteners taste sweet but have few or no calories, because taste and caloric value are independent properties of food.
The surface of a taste cell contains specialized protein molecules called taste receptors called Tlr3. When sucrose enters the active site of Tlr3, different subunits of the receptor
Split apart and this split causes ion channels in the cell membrane to open, resulting in nerve transmission that registers a sweet taste in the brain.
Artificial sweeteners taste sweet because they fit into the receptor pocket that normally
binds sucrose.

44. Chapter 10, Opener

45. Theories of
Chemical Bonding

46. Problems with Lewis Theory as a Bonding Theory
generally predicts trends in properties
does not give good numerical predictions for bond strength and bond length
A good first approximation of the bond angles in molecules, but usually cannot be used to get the actual angle
Resonance:
Cannot write a “single” correct structure for many molecules
Does not predict the correct magnetic behavior of molecules
O2 is paramagnetic, yet Lewis structure predicts it to be diamagnetic.

47. VALENCE BOND THEORY (VB) MOLECULAR ORBITAL THEORY (MO)
Theories of Bonding
VALENCE BOND THEORY (VB)
Valence electrons are localized between atoms (or are lone pairs).
Half-filled atomic orbitals overlap to form bonds.
MOLECULAR ORBITAL THEORY (MO)
Delocalization of valence electrons within molecule
Valence electrons are in orbitals (called molecular orbitals) spread over entire molecule.

48. Pi bonding occurs when:
electron orbitals of the atoms overlap sideways, resulting in electron density ABOVE and BELOW the plane of the bond’s two atomic nuclei centers.
See pi bonding with p and d orbitals, NOT with s
Sigma bonding occurs when:
electron orbitals of the atoms overlap within the plane of the bond’s two atomic nuclei centers.
Have sigma bonding with:
s to s orbitals
p to p orbitals ONLY if head to head

49. Sigma Bond (s) and Pi (P)
The electron density between two atoms
The electron density within the plane of the bond
Can be formed from hybridized or unhybridized orbitals
Every bond contains a sigma component.
Single bonds have only a sigma component.
Pi Bond
The electron density above and below the plane of the bond
Formed from unhybridized orbitals.
Pi bonds are observed in multiple bonds.
Double bonds have ONE pi component.
Triple bonds have TWO pi components.

50. Sigma Bond Formation by Orbital Overlap

51. Valence Bond Theory

52. Valence Bond Theory
It is the theory involving the overlap of atomic orbitals to form bonds.
Uses VSEPR.
The valence electrons in an atom reside in the quantum-mechanical atomic orbitals or hybrid orbitals.
The electrons must be spin paired.
A chemical bond results when these atomic orbitals interact (overlap) and there is a total of two electrons in the new molecular orbital.
When direct overlap of atomic orbital cannot or does not explain experimental finding, then hybridization of atomic orbitals of the CENTER atom occurs.
The shape of the molecule is determined by the geometry of the interacting orbitals.
The electron geometry of the molecule or ion dictates the hybridization of the atomic orbitals on the CENTER atom.
Problem:
Can NOT explain bonding with molecules or ions bringing in an odd number of valence electrons

53. Bonding in H2S
Electron configuration of each of the atom’s
valence electrons
The Lewis dot structure and distribution of hydrogen’s electrons in sulfur’s atomic orbitals coincide with the atomic orbitals of the hydrogen atoms.
The atomic orbitals on the S atom (center atom) can accommodate the 1s atomic orbitals from the hydrogen atoms without hybridization.
.. .. S H H

54. Valence Bond Theory and Hybridization
is the “mixing” of the center atom’s valence atomic orbitals
arises when the number of partially filled or empty atomic orbitals of the center atom did not predict the number of bonds or orientation of bonds between the center atom and the other bonding atom that actually occurred
Example: C = 2s22px12py12pz0 would predict two or three bonds that are 90° apart, rather than four bonds that are 109.5° apart.

55. Valence Bond Theory and Hybridization
To explain inconsistencies, it was postulated that the valence atomic orbitals of the center atom could hybridize before bonding took place.
One hybridization of C is to mix all the 2s and 2p orbitals to get four orbitals that point to the corners of a tetrahedron.

56. Bonding in a Tetrahedron Formation of Hybrid Atomic Orbitals

57. Electron Geometry (Family) Hybridization Chart

58. Predict the Bonding in Glycine Molecule

59. Predict the Bonding in Glycine Molecule

60. Predict the Bonding in Glycine Molecule

61. Predict the Bonding in Glycine Molecule

62. Predict the Bonding in Glycine Molecule

63. Hybridization and Electron Pair Geometry Relationship
Hybridization is about the center atom’s atomic orbitals undergoing a “transformation” to accommodate electrons from other atoms to form bonds.
Hybridization type corresponds to the electron geometry (family) of the compound.
*Remember: the electron geometry was determined by the number of attachments and nonbonding electrons around the center atom.
Hybridization types
Electron Geometry Hybridization Comments
linear sp triple bond
trigonal planar sp double bond
tetrahedron sp3
trigonal bipyramidal dsp3
octahedron d2sp3

64. sp2 Hybridization and Double Bonds
sp2 hybridization is seen with double bond formation.
Double bonds:
4 electrons or 2 electron pairs
Remember: not all elements can double bond
Elements that can double bond with themselves and each other are C, S, N, O, and P
Have sigma and pi components
Sigma bonds: electron density within the plane of the bond
Pi bonds: electron density above and below the plane of the bond

65. Formation of sp2 Hybrid Orbitals

66. Pi (p) Bonding in CH2O (Formaldehyde)
Formaldehyde (Formalin):
Valence electron configuration for elements in formaldehyde
Center atom in the molecule is C. Its atomic orbital needs to undergo hybridization:
The unused p orbital on the C atom contains an electron, and this p orbital overlaps the p orbital on the O atom to form a π bond.

67. Pi (p) Bonding in CH2O (Formaldehyde)
Formaldehyde (Formalin):

68. Movement Around a Double Bond
Degrees of rotation (movement) of atoms
The more electrons between two atoms in a bond, the more restrictive the rotation.
Single > Double > Triple
Bonding orbitals that form the s bond point along the internuclear axis; rotation around that bond does not require breaking the interaction between the orbitals.
Bonding orbitals that form the p bond interact above and below the internuclear axis, so rotation around the axis requires the breaking of the interaction between the orbitals.

69. sp in Bonding and Triple Bonds
sp hybridization is seen with triple bond formation.
Triple bonds:
6 electrons or 3 electron pairs
Remember: not all elements can triple bond
Elements that can triple bond with themselves and each other are C, S, N, and O
Have sigma and pi components
Sigma bonds: electron density within the plane of the bond
Two pi bonds: electron density above and below the plane of the bond
Center atom:
Uses hybrid orbitals for s bonds or lone pairs
Forms p bonds from two sets nonhybridized p orbitals
Electron and molecular geometry
linear shape
180° bond angle

70. sp in Bonding and Triple Bonds

71. sp in Bonding C2H2: Triple Bonds

72. sp3d: Hybridization for Trigonal Bipyramidal Geometries
Center atom has five electron groups (lone pairs and/or atom groups) around it.
Molecular geometry (shape)
Trigonal bipyramidal
Seesaw
T-shape
Linear
Center atom uses empty d orbitals from valence shell for hybridization.
d orbitals can be used to make p bonds.

73. sp3d: Hybridization for Trigonal Bipyramidal Geometries

74. sp3d: Hybridization AsF5

75. sp3d2: Hybridization for Octahedral Geometries
Center atom has six electron groups (lone pairs and/or atom groups) around it.
Molecular geometry (shape)
Octahedral
Square pyramidal
Square planar
Center atom uses empty d orbitals from valence shell for hybridization.
d orbitals can be used to make p bonds.

76. sp3d2: Hybridization for Octahedral Geometries

77. sp3d2 Hybridization for SF6

78. Problems with Valence Bond Theory
VB theory predicts many properties better than Lewis theory.
Bonding schemes
Bond strengths
Bond lengths
Bond rigidity
VB theory does not predict perfectly a compound’s magnetic behavior.
VB theory presumes that the electrons are localized in orbitals on the atoms in the molecule.
It doesn’t account for delocalization.

79. Molecular Orbital Theory (MO)

80. VB Theory vs. MO Theory Valance Bond
Uses VESPR and hybridization of atomic orbitals to explain bond formation
Predicts only the type of hybridization, NOT which molecules will hybridized
Can only explain bonding with compounds bringing in an even number of valence electrons
Molecular Orbital
When elements bond, a bonding orbital forms.
Bonding
Antibonding(*)
Can explain even and odd valence electron compounds
Correctly predicts bond strengths and magnetic properties of compounds
Unpaired: paramagnetic
Paired: diamagnetic
Correctly predicts whether a molecule/ion exists
Bond order

81. Molecular Orbital Theory (MO)
MO theory makes the assumption that when two elements form a bond, a new molecular orbital bonding entity is formed.
NOT the mixing of atomic orbitals, as in VB hybridization
Electrons entering MO follow Pauli’s exclusion principle (lowest to highest energy) and Hund’s rule for MO of same energy (electrons will parallel spin)
Valence electrons are delocalized.
These electrons are NOT located specifically above or below the plane of the bond.
Valence electrons are in orbitals (called molecular orbitals) spread over entire molecule.
The basis for MO theory comes from quantum mechanics.

82. Molecular Orbital Theory
Electrons in bonding MOs are stabilizing.
Lower energy than the atomic orbitals
Electrons in anti-bonding MOs are destabilizing.
Higher energy than atomic orbitals
Electron density located outside the internuclear axis
Electrons in anti-bonding orbitals cancel stability gained by electrons in bonding orbitals

83. Molecular Orbitals
When the wave functions combine constructively, the resulting molecular orbital has less energy than the original atomic orbitals and is called a bonding molecular orbital.
s, p
Most of the electron density between the nuclei
When the wave functions combine destructively, the resulting molecular orbital has more energy than the original atomic orbitals and is called an anti-bonding molecular orbital.
s*, p*
Most of the electron density outside the nuclei
Nodes between nuclei

84. Molecular Orbital Theory
1. The number of MOs formed = number of atomic orbitals used
Bonding MOs have lower energies than atomic orbitals.
Anti-bonding (*) MO is higher-level energy than atomic orbitals and bonding MOs.
Two electrons per molecular orbital
Electrons enter into to the lowest energy molecular orbital and upward.
MO Diagram

85. Illustration of Molecular Orbital: ss

86. Illustration of Molecular Orbital: sp

87. Illustration of Molecular Orbital: pp

88. Molecular Orbitals and the Property of Bond Order
Bond order is:
about strength – the greater the value, the stronger and shorter the bond between the atoms
determined by taking the difference between number of electrons in bonding and anti-bonding orbitals
Bond order = (# bonding electrons - # anti-bonding electrons) 2
Only need to consider valence electrons
May be a fraction
A bond order with a value of ZERO indicates an unstable species
Less stable than individual atoms

89. Does He2 Exist as a Molecule?
Bond order predicts whether or not a molecule or ion exists.
Bond order according to MO theory
1/2 [# e- in bonding MOs - # e- in anti- bonding MOs]
He2 bond order: ½ (2 - 2) = 0
A zero bond order indicates that He2 does not exist as a molecule.

90. Molecular Orbital Diagram of Valence Electrons for B2
B2 brings in 3 valence
electrons for each boron
atom.
Valence electrons enter
MO orbitals from lowest
energy upward.
Two electrons/orbital
Bond order
1/2 (4 - 2) = 1
Molecule is paramagnetic
Has unpaired electrons

91. Molecular Orbitals and the Property of Magnetism
Bonding theories and magnetism:
Lewis theory as well as valance bond theory do not predict a molecule’s magnetic properties.
Molecular orbital (MO) theory can predict magnetic properties.
Magnetism: results when unpaired electron occur
Materials that exhibit magnetic properties are referred to as PARAMAGNETIC.
Paramagnetic: unpaired electrons
Materials that DO NOT exhibit magnetic properties are known as DIAMAGNETIC.
Diamagnetic: all paired electrons
Example:
O2 is paramagnetic.

92. Molecular Orbitals and the Property of Magnetism in O2
O2 is paramagnetic.
MO diagram for oxygen

93. Predict the more stable species:
___ * 2p *2p ___ ___ *2p 2p ___ ___ 2p ___  2p ___ * 2s ___  2s
___ * 2p *2p ___ ___ *2p 2p ___ ___ 2p ___  2p ___ * 2s ___  2s
___ * 2p *2p ___ ___ *2p 2p ___ ___ 2p ___  2p ___ * 2s ___  2s
Bond order = 2
Bond order = 2 1/2
Bond order = 1 1/2

94. MO Diagrams of 2nd Row Nonmetals