1. Lewis Theory of Bonding
Chemical Bonding
Lewis Theory of Bonding

2. Types of bonding conditions between elements
Low Electronegativity and low Ionization energy (Metals)
High electronegativity and High Ionization energy (Non-metals)
Metallic bonding
Ionic bonding (transferring of electrons between atoms)
Ionic bonding
Covalent bonding (sharing of electrons between atoms)

3. Electronegativity

4. Lewis Diagrams or Structures
A convention developed to “show” the relationship between atoms when they form bonds.
Why is it necessary?
predict where the electrons are in a molecule
needed to predict the shape of a molecule

5. Lewis Diagrams for Ionic Compounds
Identify the number of valence shell electrons and determine the charge on the ion using the “stable octet rule”.
Write the elemental symbol, place dots to represent the electrons in the valence shell, enclose in square brackets and write the ionic charge as a superscript.
[Na]+ or [ Cl ]-

6. Lewis Diagrams for Covalent Compounds
Draw the Lewis Diagram for nitrogen trifluoride (NF3). Step 1. Count the valence electrons N = 5 F = ( 7) = 26 valence electrons

7. Lewis Diagrams for Covalent Compounds
Step 2. Write a skeletal structure. Use the least electronegative atom in the centre Electronegativity: N = 3.0 & F = 4.0 F N
= a pair of e-
(a single bond)

8. Lewis Diagrams for Covalent Compounds
Step 3. Complete the octets for each terminal atom (except H) F N : ..

9. Lewis Diagrams for Covalent Compounds
Step 4. Assign any additional electrons as lone pairs on the central atom .. F N :

10. Lewis Diagrams for Covalent Compounds
Example 2. COCl2 (24 electrons) .. .. : : Cl C Cl .. .. : : O ..

11. Lewis Diagrams for Covalent Compounds
Step 5. Make multiple bonds where necessary to complete the octets. .. C Cl O : C Cl O : ..

12. Lewis Diagrams for Covalent Compounds
Example 3. Chlorate ion, ClO3- ((1 x 7) + (3 x 6) + 1) = 26 - .. .. .. : : O Cl O .. .. : : O ..

13. Co-ordinate covalent bonds
In some covalent compounds, the bonds between atoms occur because one atom has donated both electrons to the covalent bond. This is called a coordinate covalent bond. N H +
N : H H+ +
Nitrogen supplies the two lone pair electrons to this N-H bond. The H+ ion has no electrons.

14. Co-ordinate covalent bonds
To determine the number of coordinate covalent bonds – subtract the bonding capacity (lone valence electrons) from the number of bonds the atom has. N H +
Nitrogen Bonds  4 Bonding capacity 3 Coordinate bonds  4-3=1

15. In some compounds the SCH3U guidelines may not “work”.
On occasion, both elements have the same electronegativity or there may be two or more possible Lewis Structures.
e.g. CS2
(both electronegativities = 2.5)
is it S=C=S or C=S=S ?

16. Formal Charge
In such situations, one determines the Formal Charge. The option with the lowest formal charge has the most stable and viable structure.
The Formal Charge for an atom is the number of valence electrons in the free neutral atom minus the number of valence electrons assigned to the atom in the Lewis structure.

17. Formal Charge
Formal Charge = (# valence electrons)-(# of bonds)-(# of unshared e-)

18. C=S=S : ..
C S S Valence electrons Electrons assigned Formal Charge

19. S=C=S .. :
S C S Valence electrons Electrons assigned Formal Charge 0 0 0

20. Resonance Structures
In some structures the Lewis structure does not represent the true structure of the compound.
Bond order is the number of shared pairs of electrons between two atoms. (i.e. – the number of bonds between two atoms)
As the bond order increases. . .
The length of the bond decreases.
The energy associated with breaking the bond increases.

21. Bond order
The number of shared electrons in a bond affects its length and energy.
Bond Type
Bond Order
Bond Length (pm) (10-12m)
Bond Energy (kJ/mol)
C-O 1
143
351
C=O 2
121
745
C-C
154
348
C=C
134
615
C C 3
120
812
C-N
276
C=N
138
C N
116
891

22. Resonance structure C H O -1
CHO2- is a polyatomic ion with the following Lewis structure.
The C-O bond lengths are experimentally determined to be between C-O and C=O. The bond order is neither 1 or 2, but considered to be somewhere in between (i.e.-1.5).
The Lewis structure does not support the experimental data.

23. Resonance structure
The actual structure is a resonance hybrid of the two resonance structures. C H O -1
The resonance structure does not “flip-flop” back and forth between the two. It is a hybrid form of the two.

24. Resonance structure The bond order for NO3- is, (1+1+2)/3=1.33.
The resonance structure is N O -1

25. Benzene has a bond order of 1.5. (1+2+1+2+1+2)/6=1.5C H C H

26. Nature of the chemical bond
The work on quantum theory in conjunction with the success of Lewis structures resulted in the inevitable connections between the two areas of study.
Linus Pauling , a friend of Gilbert Lewis, connected the two with the valence bond theory.

27. Valence Bond Theory
The half filled orbital of one atom overlaps a half filled orbital of a second atom to form a full orbital with two electrons spinning in opposite directions.
The bonding atoms arrange themselves in order to maximize the overlap of the half-filled orbitals.
Maximum overlapping of the orbitals creates a bonding orbital with a lower energy and increased stability.

28. Valence Bond Theory
The combination of electrostatic repulsion and opposing magnetic fields (due to the electron’s spin) creates the stability associated with a bonding orbital.
Electron Spin
Negative charge
Electrostatic repulsion
Negative charge
Electron Spin
“North” magnetic field
Magnetic attraction
“South” magnetic field

29. Valence Bond Theory
During this process, two atoms approach each other and allow their half filled orbitals to overlap and form the stability of a filled bonding orbital.

30. Hybrid Orbitals
In some situations a more advantageous bonding scenario can be established by promoting electrons from a full orbital to a similar empty orbital to create two half filled orbitals that are available for bonding.
The resulting orbital is a mixture of the two original orbitals and is called a hybrid orbital.

31. Hybrid Orbitals
Electrons are promoted from the full “s-orbital” into an empty “p-orbital” which results in hybrid orbitals that have one electron per orbital and characteristics unique to the newly formed orbitals.
The new hybrid orbitals are free to become involved in bonds by overlapping with other half filled valence orbitals.

32. Hybridization C (z = 6) C (z = 6) C (z = 6) C (z = 6) C (z = 6)1s 2s 2p
C (z = 6) 1s
sp3
C (z = 6)
sp3 hybridization 1s 2s 2p
C (z = 6) 1s
sp2
C (z = 6)
sp2 hybridization p 1s 2s 2p
C (z = 6) 1s sp
C (z = 6)
sp hybridization p 1s 2s 2p
C (z = 6)

33. sp2 Hybrid atomic orbitals

34. sp3 Hybrid atomic orbitals

35. Double & Triple Bonds
Sigma (s) bonds – The “end-to-end” overlapping of half filled orbitals to make a full bonding orbital of lower energy level (i.e. – more stable) They occur between “s”, “p” and hybrid orbitals (“sp”,“sp2” & “sp3”) to make single covalent bonds.

36. Double & Triple Bonds
Pi (p) bonds – The “side-to-side” overlapping of half filled “p” orbitals to make more stable filled bonding orbitals.

37. Double & Triple Bonds
A combination of “p” and “s” bonds makes double and triple bonds.
Single bonds Sigma only (s)
Double bonds 1 Sigma (s) and 1 Pi (p)

38. Triple bonds 1 Sigma (s) and 2 Pi (p)
Double & Triple Bonds
A combination of “p” and “s” bonds makes double and triple bonds.
Sigma (s ) bond
Pi (p ) bond
Pi (p ) bond
Triple bonds 1 Sigma (s) and 2 Pi (p)

39. Molecular shapes Valence shell electron pair repulsion (VSEPR) theory
Only valence shell electrons of the central atom are important in the molecular shape.
Valence shell electrons will repel to maximize the distance between the pairs
Bond pairs and lone pairs behave similarly

40. Molecular shapes

47. 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-