1. Molecular Shape and Theory of Chemical Bonding
Shapes of Molecules and Polyatomic Ions
Polar and Nonpolar Molecules
Bonding Theory
Molecular Orbital Method
Delocalized Electrons
Band Theory of Bonding in Solids

2. Shapes of molecules and polyatomic ions
Molecules and polyatomic ions are not all ‘flat’ structures.
Many have a three dimensional arrangement that helps account for their various chemical and physical properties.
Several models are used to help predict and describe the geometries for these species.
One model is called the Valence Shell Electron Pair Repulsion model (VSEPR)

3. VSEPR model
According to this model, for main group elements, electron pairs will be as far apart from each other as possible.
This occurs in three dimensional space.
Both bonded and unshared pairs will occupy space with unshared pairs taking up more space.
The geometry is based on the total number of electron pairs - total coordination number.

4. VSEPR shapes 2 2 0 AB2 Linear 3 3 0 AB3 Trigonal planar 2 1 AB2 Bent
Coordination Electron pairs General
Number Bonding Unshared Formula Shape
AB Linear
AB Trigonal planar
AB Bent
AB Tetrahedral
AB Trigonal pyramidal
AB Bent
AB Linear

5. Molecular geometry Molecules have specific shapes.
Determined by the number of electron pairs around the central species
Bonded and unshared pairs count.
Multiple bonds are treated as a single bond for geometry.
Geometry affects factors like polarity and solubility.

6. Some common geometries
e- pairs around
Shape central atom Example
Linear BeH2
Trigonal planar BF3
Tetrahedral CH4
Trigonal pyramidal NH3
Bent H2O

7. Linear - CO2

8. Trigonal planar, BCl3

9. Bent, H2O

10. Pyramidal, NH3

11. Tetrahedral, CH4

12. Molecular geometries based on tetrahedral
Pyramidal C
Tetrahedral
Bent O
Bent and pyramidal are
actually tetrahedral but
some of the electron
pairs are not bonded.

13. Other geometries. Other shapes are also observed.
Five bonds or lone electron pairs
Trigonal bipyramidal
Seesaw
T-shaped
Linear
Six bonds or lone electron pairs
Octahedral
Square pyramidal
Square planar

14. VSEPR shapes 5 5 0 AB5 Trigonal bipyramidal 4 1 AB4 Seesaw
Coordination Electron pairs General
Number Bonding Unshared Formula Shape
AB Trigonal
bipyramidal
AB Seesaw
AB T-shaped
AB Linear
AB Octahedral
AB Square
pyramidal
AB Square Planar

15. Trigonal bipyramidal

16. Square planar

17. Octahedral

18. Molecular geometry H C As molecules get larger, the rules regarding
molecular geometry still hold.

19. Ethane

20. Geometry and polar molecules
For a molecule to be polar
- must have polar bonds
- must have the proper geometry
CH4 non-polar
CH3Cl polar
CH2Cl2 polar
CHCl3 polar
CCl4 non-polar
WHY?

21. Polar and nonpolar molecules
Polarity is an important property of molecules.
It affects physical properties such as melting point, boiling point and solubility.
Chemical properties also depend on polarity.
Dipole moment, m, is a quantitative measure of the polarity of a molecule.

22. Dipole moment This property can be measured by placing molecules
in an electrical field. Polar molecules will align when
The field is on. Nonpolar molecules will not. + - + -

23. Polar and nonpolar molecules
Most bonds between atoms of dissimilar elements in a molecule are polar. That does not mean that the molecule will be polar.
Electronegativities:
Oxygen = 3.5
Carbon = 2.5
Difference 1.0
(polar bond)
O = C = O
The electronegativity values
Show that the C-O bond would be polar with electrons
Being pulled towards the oxygens. However, due to
The geometry, the pull happens in equal and opposite
directions.

24. Polar and nonpolar molecules
For a molecule to be polar, the effects of bond polaritymust not cancel out.
One way is to have a geometry that is not symmetrical like in water.
H H O ..
Electronegativity
difference = 1.3
Here, the effects of the polar bonds do not
canceled so the molecule is polar.

25. Polar and nonpolar molecules
A molecule is nonpolar if the central atom is symmetrically substituted by identical atoms.
CO2, CH4 , CCl4
A molecule will be polar if the geometry is not symmetrical.
H2O, NH3, CH2Cl2
The degree of polarity is a function of the number and type of polar bonds as well as the geometry.

26. Bonding theory
Two methods of approximation are used to describe bonding between atoms.
Valence bond method
Bonds are assumed to be formed by overlap of atomic orbitals
Molecular orbital method
When atoms form compounds, their orbitals combine to form new orbitals - molecular orbitals.

27. Valence bond method
According to this model, the H-H bond forms as a result of the overlap of the 1s orbitals from each atom.
74 pm

28. Valence bond method
Hybrid orbitals are need to account for the geometry that we observe for many molecules.
Example - Carbon
Outer electron configuration of 2s2 2px1 2py1
We know that carbon will form2 four equivalent bonds - CH4, CH2Cl2 , CCl4.
The electron configuration appears to indicate that only two bonds would form and they would be at right angles -- not tetrahedral angles.

29. Hybridization
To explain why carbon forms four identical single bonds, we assume the the original orbitals will blend together.
Unhybridized Hybridized
energy 2s 2p
2sp3

30. Hybridization
In the case of a carbon that has 4 single bonds, all of the orbitals are hybrids.
sp3
25% s and 75% p character 1
+ 3 4
s p sp3

31. Ethane, CH3CH3  bond sp3 hybrids s bond - formed by an endwise
1s orbital
of H
 bond
sp3
hybrids
s bond - formed
by an endwise
(head-on) overlap.
Molecules are
able to rotate
around single
bonds.

32. Ethane , CH3CH3

33. Rotation of single bond

34. Rotation of single bond

35. sp2 hybrid orbitals
To account for double bonds, a second type of hybrid orbital must be pictured. An sp2 hybrid is produced by combining one s and 2 p orbitals. One p orbital remains. 2p 2p
energy
2sp2 2s
Unhybridized Hybridized

36. sp2 hybrid orbitals
The unhybridized p orbitals are able to overlap, resulting in the formation of a second bond - p bond.
A p bond is a
sideways overlap
that occurs both
above and below the
plane of the molecule
Parts of the molecule
are no longer able to
rotate about the bond. C

37. Ethene

38. Bonding in ethene  bond  bond 1s orbital p overlap sp2 hybrids sp2

39. Bonding in ethene

40. sp hybrid orbital
Forming a triple bond is also possible. This requires that two p orbitals remain unhybridized. 2p 2p
energy
2sp 2s
Unhybridized Hybridized

41. sp hybrid orbital
Now two p orbitals are available to
form p bonds.

42. Ethyne

43. Bonding in ethyne
sp hybrid
p overlaps

44. Bonding in ethyne

45. Other hybrid orbitals
d orbitals can also be involved in the formation of hybrid orbitals.
Hybrid Shape
sp Linear
sp2 Trigonal planar
sp3 Tetrahedral
sp3d Trigonal bipyramidal
sp3d2 Octahedral

46. Molecular Orbital Method
When atomic orbitals combine to form molecular orbitals, the number of molecular orbitals formed must equal the number of atomic orbitals mathematically combined.
Example - H2
Two 1s orbitals will combine forming two molecular orbitals. The overall energy of the new orbitals is the same as the original two 1s. However, they will be at different energies.

47. H2 molecular orbital diagram1s
s1s H2
s*1s
energy
Orbital shapes

48. Molecular orbitals
When two atomic orbitals combine, three types of molecular orbitals are produced.
Bonding orbital - s or p
The energy is lower than the atomic orbitals and the electron density overlaps.
Antibonding orbital - s* or p*
The energy is higher than the atomic orbitals and the electron density does not overlap.
Nonbonding - n
Electron pairs not involved in bonding.

49. Homonuclear diatomic molecules
These molecules are simple diatomics where both atoms are of the same element.
Energy diagrams for these types of molecules are similar to the one for H2.
We can develop energy diagrams for a range of molecules or possible molecules to see if they bond and how.

50. MO diagram of helium s*1s s1s If we develop a diagram for helium
we see that both
a bonding and
antibonding
orbital will be
filled.
The result is that
it is no more stable
than the unbonded
form -- it will not
bond He 1s
s1s
He2
s*1s
energy

51. Molecular orbital bonding
For a molecule to be stable, you must have more electrons in bonding orbitals than in antibonding orbitals.
The bonded form will be at a lower energy so will be more stable.
Bonding and antibonding orbitals for both s and p bonds must be considered.
Lets look at the MO diagram for O2.

52. MO diagram for O2 1s 2s 2px 2py 2pz s*2pz s2pz s2s s*2s s*1s s1s p2px
p2py
p*2px
p*2py

53. MO diagram for O2 Each oxygen atom has 8 electrons for a total of 16.
We can now place 16 electrons into the MO diagram and see what happens.
Remember, don’t pair electrons unless you need to and fill a lower energy orbital before proceeding to the next higher one.
O2 will form if we have more bonding than antibonding electrons.

54. MO diagram for O2 1s 2s 2px 2py 2pz s*2pz s2pz s2s s*2s s*1s s1s p2px
p2py
p*2px
p*2py

55. Heteronuclear diatomic molecules
Molecular orbital diagrams become more complex when bonding between two nonidentical atoms is considered.
The atomic energy levels are not the same and there are differing numbers of electrons.
A simple example is NO where the orbitals are similar but not identical.

56. MO diagram for NO 1s 2s 2px 2py 2pz s*2px s2px s2s s*2s s*1s s1s p2py
p2pz
p*2py
p*2pz N NO O

57. Delocalized electrons
MO diagrams for polyatomic species are often simplified by assuming that all s and some p orbitals are localized -- shared between two specific atoms.
Resonance structures require that electrons in some p orbitals be pictured as delocalized.
Delocalized - free to move around three or more atoms.

58. Delocalized electrons
Benzene is a good example of delocalized electrons.
We know that the bonding between carbons has an order of 1.5 and that all of the bonds are equal. =

59. Aromatic hydrocarbons
p orbitals overlap
sidewise all around
the ring. No localized
double bonds.

60. Band theory of bonding in solids
This is an extension of delocalized orbitals.
Each atom interacts with all of the others in the crystal, resulting in an enormous number of ‘molecular orbitals.’ 3s
9 Na 3s
9 Na

61. Band theory of bonding in solids
A group of very closely spaced energy levels.
Energy gap
The difference in energy between the bonding and antibonding orbitals.
Forbidden bands
A ‘space’ that separates bands.

62. Band theory of bonding in solids
The s and p bands of
Group II (2)
metals overlap. p
Energy s
Internuclear distance

63. Band theory of bonding in solids
Conductor
A material with a partially filled energy band.
Insulator
The highest occupied band is filled or almost completely filled. The forbidden band just above the highest filled is wide.
Semiconductor
The gap between the highest filled band and the next higher permitted band is relatively narrow.

64. Band theory of bonding in solids
Empty
Energy
Forbidden, wide
Filled
insulator
Empty
Energy
Forbidden, narrow
Filled
semiconductor
Energy No
Forbidden
conductor