1. The Shapes of Molecules
The shape of a molecule has an important bearing on its reactivity and behavior.
The shape of a molecule depends a number of factors. These include:
Atoms forming the bonds
Bond distance
Bond angles 1

2. Valence Shell Electron Pair Repulsion
Valence Shell Electron Pair Repulsion (VSEPR) theory can be used to predict the geometric shapes of molecules.
VSEPR is revolves around the principle that electrons repel each other.
One can predict the shape of a molecule by finding a pattern where electron pairs are as far from each other as possible. 2

3. Bonding Electrons and Lone Pairs
In a molecule some of the valence electrons are shared between atoms to form covalent bonds. These are called bonding electrons.
Other valence electrons may not be shared with other atoms. These are called non-bonding electrons or they are often referred to as lone pairs. 3

4. VSEPR
In all covalent molecules electrons will tend to stay as far away from each other as possible
The shape of a molecule therefore depends on:
the number of regions of electron density it has on its central atom,
whether these are bonding or non-bonding electrons. 4

5. Hybridization - The Blending of Orbitals+ =
Poodle +
Cocker Spaniel =
Cockapoo + =
s orbital +
p orbital =
sp orbital

6. What Proof Exists for Hybridization?
We have studied electron configuration notation and
the sharing of electrons in the formation of covalent
bonds.
Lets look at a molecule of methane, CH4.
Methane is a simple natural gas. Its molecule has a
carbon atom at the center with four hydrogen atoms covalently bonded around it.

7. Carbon ground state configuration
What is the expected orbital notation of carbon
in its ground state?
Can you see a problem with this?
(Hint: How many unpaired electrons does this
carbon atom have available for bonding?)

8. Carbon’s Bonding Problem
You should conclude that carbon only has TWO
electrons available for bonding. That is not not enough!
How does carbon overcome this problem so that
it may form four bonds?

9. Carbon’s Empty Orbital
The first thought that chemists had was that carbon promotes one of its 2s electrons…
…to the empty 2p orbital.

10. However, they quickly recognized a problem with such
an arrangement…
Three of the carbon-hydrogen bonds would involve
an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.

11. This would mean that three of the bonds in a methane
molecule would be identical, because they would involve
electron pairs of equal energy.
But what about the fourth bond…?

12. The fourth bond is between a 2s electron from the
carbon and the lone 1s hydrogen electron.
Such a bond would have slightly less energy than the other bonds in a methane molecule.

13. This bond would be slightly different in character than
the other three bonds in methane.
This difference would be measurable to a chemist
by determining the bond length and bond energy.
But is this what they observe?

14. The simple answer is, “No”.
Measurements show that
all four bonds in methane
are equal. Thus, we need
a new explanation for the
bonding in methane.
Chemists have proposed an explanation – they call it
Hybridization.
Hybridization is the combining of two or more orbitals
of nearly equal energy within the same atom into
orbitals of equal energy.

15. In the case of methane, they call the hybridization
sp3, meaning that an s orbital is combined with three
p orbitals to create four equal hybrid orbitals.
These new orbitals have slightly MORE energy than
the 2s orbital…
… and slightly LESS energy than the 2p orbitals.

16. sp3 Hybrid Orbitals
Here is another way to look at the sp3 hybridization
and energy profile…

17. sp Hybrid Orbitals While sp3 is the hybridization observed in methane,
there are other types of hybridization that atoms
undergo.
These include sp hybridization, in which one s
orbital combines with a single p orbital.
This produces two hybrid orbitals, while leaving two normal p orbitals

18. sp2 Hybrid Orbitals
Another hybrid is the sp2, which combines two orbitals from a p sublevel with one orbital from an s sublevel.
One p orbital remains unchanged.

19. Hybridization Involving “d” Orbitals
Beginning with elements in the third row, “d” orbitals may also hybridize
sp3d = five hybrid orbitals of equal energy
sp3d 2 = six hybrid orbitals of equal energy

20. VSEPR Predicting Shapes

21. VSEPR: Predicting the shape
Once the dot structure has been established, the shape of the molecule will follow one of basic shapes depending on:
The number of regions of electron density around the central atom
The number of regions of electron density that are occupied by bonding electrons 21

22. VSEPR: Predicting the shape
The number of regions of electron density around the central atom determines the electron geometry
The number of regions of electron density that are occupied by bonding electrons and hence other atoms determines the actual shape 22

23. Linear Molecules
Linear molecules have only two regions of electron density. 23

24. Angular or Bent
Angular or bent molecules have at least 3 regions of electron density, but only two are occupied 24

25. Triangular Plane
Triangular planar molecules have three regions of electron density. All are occupied by other atoms 25

26. Tetrahedron All are occupied by other atoms
Tetrahedral molecules have four regions of electron density.
All are occupied by other atoms 26

27. Trigonal Bipyramid
A few molecules have expanded valence shells around the central atom. Hence there are five pairs of valence electrons. The structure of such molecules with five pairs around one is called trigonal bipyramid. 27

28. Octahedron
A few molecules have valence shells around the central atom that are expanded to as many as six pairs or twelve electrons. These shapes are known as octahedrons 28

29. VSEPR – Valence Shell Electron Pair Repulsion
X + E
Overall Structure
Forms 2
Linear
AX2 3
Trigonal Planar
AX3, AX2E 4
Tetrahedral
AX4, AX3E, AX2E2 5
Trigonal bipyramidal
AX5, AX4E, AX3E2, AX2E3 6
Octahedral
AX6, AX5E, AX4E2
A = central atom
X = atoms bonded to A
E = nonbonding electron pairs on A

30. VSEPR: Linear
linear
AX2
CO2

31. VSEPR: Trigonal Planar
AX3
BF3
bent
AX2E
SnCl2

32. VSEPR: Tetrahedral AX4 CCl4 AX3E PCl3 AX2E2 Cl2O tetrahedral
Trigonal pyramidal
AX3E
PCl3
bent
AX2E2
Cl2O

33. VSEPR: Trigonal Bi-pyramidal
Trigonal planar
AX5
PCl5
Trigonal pyramidal
AX4E
SF4
T-shaped
AX3E2
ClF3
linear
AX2E3
I3-

34. VSEPR: Octahedral AX6 SF6 AX5E BrF5 AX4E2 ICl4- octahedral
Square pyramidal
AX5E
BrF5
Square planar
AX4E2
ICl4-

35. Sigma and Pi Bonds
Sigma () bonds exist in the region directly between two bonded atoms.
Pi () bonds exist in the region above and below a line drawn between two bonded atoms.
Single bond
1 sigma bond
Double Bond
1 sigma, 1 pi bond
Triple Bond
1 sigma, 2 pi bonds
Visual of how this happens…

36. Bond Length Diagram

37. Atomic orbitals overlap
When covalent bonds are formed, the orbitals overlap to create a new molecular orbital.
Simply put, this happens because the new orbital is of lower energy.
There are two main types.

38. The sigma bond
When two atomic orbitals overlap along the bond axis (between the two nuclei)
(single bonds)

39. The pi bond
When two p orbitals overlap sideways in the space above and below the bond axis
(double and triple bonds)

40. Identifying sigma and pi orbitals

41. Sp3 hybridization
When four single bonds are formed, the s and p orbitals blend creating four new orbitals
Ex. CH4

42. Sp2 hybridization
When a double bond is formed, three equal orbitals are produced
Ex. C2H4

43. Sp hybridization
Triple bonds = two equal orbitals
Ex. C2H2

44. Hybridization to predict molecular shape

45. Sigma and Pi Bonds Single Bonds
Ethane

46. Sigma and Pi Bonds: Double bonds
Ethene
1  bond

47. Sigma and Pi Bonds Triple Bonds
Ethyne
1  bond
1  bond

48. The De-Localized Electron Model
Pi bonds () contribute to the delocalized model of electrons in bonding, and help explain resonance
Electron density from  bonds can be distributed symmetrically all around the ring, above and below the plane.