1. Chemical Bonding

2. Chemical Bonds, Lewis Symbols, and the Octet Rule
Chemical bond: attractive force holding two or more atoms together.
Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals.
Ionic bond results from the transfer of electrons from a metal to a nonmetal.
Metallic bond: attractive force holding pure metals together.

3. Ionic Bonding

4. Electron Configurations of Ions and their Lewis Structures

5. Covalent Bonding

6. Energy Potential Diagram for Covalent Bonds!

7. Chemical Bonds Bond Type Single Double Triple # of e’s 2 4 6 Notation— = 
Bond order 1 3
Bond strength
Increases from Single to Triple
Bond length
Decreases from Single to Triple

9. Strengths of Covalent Bonds

11. Chemical Bonds, Lewis Symbols, and the Octet Rule

12. Chemical Bonds, Lewis Symbols, and the Octet Rule
All noble gases except He has an s2p6 configuration.
Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs).
Caution: there are many exceptions to the octet rule.

13. Bond Polarity and Electronegativity
Electronegativity: The ability of one atoms in a molecule to attract electrons to itself.
Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F).
Electronegativity increases
across a period and
down a group.

14. Figure 8.6: Electronegativities of Elements
Electronegativity

15. Difference in Electronegativities
0 to 0.5 Nonpolar covalent
0.5 to 1.7 Polar covalent
>1.7 Ionic

16. Bond Polarity and Electronegativity
Figure 8.7: Electronegativity and Bond Polarity
There is no sharp distinction between bonding types.
The positive end (or pole) in a polar bond is represented + and the negative pole -.
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17. Drawing Lewis Structures Follow Step by Step Method
Total all valence electrons. [Consider Charge]
Use Gr # on PT. ex Gr1=1v.e, Gr2=2ve etc
Write symbols for the atoms and guess skeleton structure [ define a central atom ].
Place a pair of electrons in each bond.
Complete octets of surrounding atoms. [ H = 2 only ]
Place leftover electrons in pairs on the central atom.
If there are not enough electrons to give the central atom an octet, look for multiple bonds by transferring electrons until each atom has eight electrons around it.

18. Lewis Dot Diagrams for Individual Elements

19. Exceptions to the Octet Rule
Central Atoms Having Less than an Octet
Relatively rare.
Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A.
Most typical example is BF3.
Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.

20. Exceptions to the Octet Rule
Central Atoms Having More than an Octet
This is the largest class of exceptions.
Atoms from the 3rd period onwards can accommodate more than an octet.
Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.
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21. Drawing Lewis Structures
Formal Charge
Consider:
For C:
There are 4 valence electrons (from periodic table).
In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure.
Formal charge: = -1.

22. Drawing Lewis Structures
Formal Charge
Consider:
For N:
There are 5 valence electrons.
In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure.
Formal charge = = 0.
We write:

23. Formal Charge facts to keep in mind:
Formal charges need to cancel out on a neutral compound. Example SO3
Formal charges equal (when added) the total charge of the compound. Example NO3-1

24. Molecular Shapes: VSEPR
There are five fundamental geometries for molecular shape:

25. Molecular Shapes – 3D Notations
VSEPR (Ballons)-Movie Clip

26. Figure 9.3
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27. Summary of VSEPR Molecular Shapes
e-pairs
Notation
Name of VSEPR shape
Examples 2
AX2
Linear
HgCl2 , ZnI2 , CS2 , CO2 3
AX3
Trigonal planar
BF3 , GaI3
AX2E
Non-linear (Bent)
SO2 , SnCl2 4
AX4
Tetrahedral
CCl4 , CH4 , BF4-
AX3E
(Trigonal) Pyramidal
NH3 , OH3-
AX2E2
Non-Linear (Bent)
H2O , SeCl2 5
AX5
Trigonal bipyramidal
PCl5 , PF5
AX4E
Distorted tetrahedral
(see-sawed)
TeCl4 , SF4
AX3E2
T-Shaped
ClF3 , BrF3
AX2E3
I3- , ICl2- 6
AX6
Octahedral
SF6 , PF6-
AX5E
Square Pyramidal
IF5 , BrF5
AX4E2
Square Planar
ICl4- , BrF4-

28. Examples: VSEPR Molecular Shapes - I
# electron pairs on Central Atom A
Notation
Example
Lewis
VSEPR & Name of Shape 2
AX2
2 bp on A 3
AX3
3 bp on A
AX2E
2 bp and 1 lp on A

29. Examples: VSEPR Molecular Shapes – I – F08

30. Examples: VSEPR Molecular Shapes - II
# electron pairs on Central Atom A
Notation
Example
Lewis
VSEPR & Name of Shape 4
AX4
4 bp on A
AX3E
3 bp and 1 lp on A
AX2E2
2 bp and 2 lp on A

31. Examples: VSEPR Molecular Shapes – II – F08

32. Examples: VSEPR Molecular Shapes - III
# electron pairs on Central Atom A
Notation
Example
Lewis
VSEPR & Name of Shape 5
AX5
5 bp on A
AX4E
4 bp and 1 lp on A
AX3E2
3 bp and 2 lp on A
AX2E3
2 bp and 3 lp on A

33. Examples: VSEPR Molecular Shapes – III – F08
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34. Examples: VSEPR Molecular Shapes - IV
# electron pairs on Central Atom A
Notation
Example
Lewis
VSEPR & Name of Shape 6
AX6
6 bp on A
AX5E
5 bp and 1 lp on A
AX4E2
4 bp and 2 lp on A

35. VSEPR Model The Effect of Nonbonding Electrons
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 increases
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36. VSEPR Model Figure 9.10: Shapes of Larger Molecules
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Figure 9.10: Shapes of Larger Molecules
In acetic acid, CH3COOH, there are three central atoms.

37. Figure 8.10: Drawing Lewis Structures
Resonance Structures

38. HyperChem
Figure 9.12

39. Figure 9.11: Molecular Shape and Molecular Polarity
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40. Figure 9.13: Molecular Shape and Molecular Polarity
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41. Covalent Bonding and Orbital Overlap
Lewis structures and VSEPR do not explain why a bond forms.
How do we account for shape in terms of quantum mechanics?
What are the orbitals that are involved in bonding?
We use Valence Bond Theory:
Bonds form when orbitals on atoms overlap.
There are two electrons of opposite spin in the orbital overlap.

42. Figure 9.14: Covalent Bonding and Orbital Overlap
Optional Topic

45. VSEPR Model (Figure 9.6) To determine the electron pair 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.

46. VSEPR Model

49. Chemical Bonding
Lewis
VSEPR shapes
AXE notation
Polarity