1. Chemical Bonding I: Basic Concepts
Chapter 9

2. Chapter 9 Topics Bonding: Orbitals Lewis dot structures Ionic bond
Covalent bond
Electronegativity and bond polarity
Writing Lewis structure (Octet Rule)
Exception to the octet rule
Resonance
Formal charge
Bond energy and enthalpy
Bonding: Orbitals

4. Chemical Bonds
(chemical bond in which electron completely transfer form the metal to the nonmetal)
1) Ionic Bond:
Metal nonmetal Ionic compound
v. Low e affinity
High e affinity
v. Low electronegativity
High electronegativity
Tend to lose e
Tend to gain e
Li (1s2 2s1) Li+ (1s2) + e He
F (1s2 2s22p5) + e F- (1s22s22p6) Ne Li + F
Li+ -
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5. (chemical bond in which two or more electrons are shared by two atoms)
2) Covalent Bond
nonmetal nonmetal Covalent compound F +
7e-
8e-
nonpolar bond and nonpolar molecule
H Cl H-Cl
polar bond and polar molecule
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6. (the attraction that an atom has for the electrons in a bond)
Electronegativity
(the attraction that an atom has for the
electrons in a bond)
This is different than electron affinity which is about isolated atom
Resultant = 0
Resultant =
H H
H F + -
Same electronegativity
(nonpolar bond)
For diatomic it is also
nonpolar compound
different electronegativity
(polar bond)
For diatomic it is also polar compound
Any diatomic molecule formed from two element of different
electronegativity is a polar molecule and have a dipole moment

7. Dr. Ali Bumajdad

8. Dr. Ali Bumajdad

9. Bond Polarity and Dipole moments
The grater the difference between the electronegativity of the
atoms in a compound, the greater the polarity of the compound
H Cl
=0.9
N O
=0.5
C O
=1.0
H F
=1.9 3 4 2 1
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10. Classification of bonds by difference in electronegativity
Bond Type
Covalent
 2
Ionic
0 < and <2
Polar Covalent
Increasing difference in electronegativity
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
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11. Q) Classify the following bonds as ionic, polar covalent,
or covalent: The bond in CsCl; the bond in H2S; and
the NN bond in H2NNH2.
Cs – 0.7
Cl – 3.0
3.0 – 0.7 = 2.3
Ionic
H – 2.1
S – 2.5
2.5 – 2.1 = 0.4
Polar Covalent
N – 3.0
N – 3.0
3.0 – 3.0 = 0
Covalent
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12. H-H < H-S < H-Cl < O-H < F-H
Sa. Ex. : Order according to polarity:
H-H, O-H, H-Cl, H-S, F-H
H-H < H-S < H-Cl < O-H < F-H
Covalent bond
Polar covalent bond
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13. Polar bonds but nonpolar molecule Polar bonds and polar molecule
O = C = O O H
Polar bonds but nonpolar molecule
Polar bonds and polar molecule
Sa. Ex. : Which ones have a dipole moment:
HCl, Cl2, SO3, CH4, H2S
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14. Lewis Structures (Octet Rule)
Octet rule: An atom tend to lose or gain or share electrons until
its outer shell (valence shell) contain 8 electrons (nobel gas
configuration)
Lewis Symbol:
Help us knowing number of bonds but not number of
unpaired electrons
3 bonds
3 v.e.
2 bonds
6 v.e. B O

15. Drawing Lewis Structures:
A) Deciding the central atom:-
In general, the central atom is the first atom in the binary molecules
e.g. : CO2, CO32-, NH3, NO2, NO3-, SO3, SO42-
exceptions are : H2O and H2S
B) Distributing the valence electrons:-
1) Count all the valence electrons
2) Place one pair of electrons in each bond
3) Complete the octets for the non-central atoms
4) Place the remaining electrons on the central atom.
5) If the central atom still has less than an octet, use multiple bonds

16. Sa. Ex. : Write Lewis structure for :
HF b) N2
NH3 d) CH4
CF4 e) NO+
f) CO2 f) NO3-
g) CO g) H2S

17. central atom with principal quantum number n > 2
Exception of the octet rule
The Incomplete Octet
1) BeCl2
2) BF3 or BCl3
central atom with principal quantum number n > 2
3) SF6
4) PCl5
Odd-Electron Molecules
5) NO N O

18. Sa. Ex. : Write Lewis structure for :
ClF3 b) XeO3
RnCl2 d) BeCl2
e) ICl4-

19. Resonance
A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure because the bond length measurement show that the actual bond length is similar.
e.g. O + - O + - O C - O C - O C -
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20. Q) What are the resonance structure of NO2- and NO3-

21. (Charge calculated from an atom in the Lewis structure)
Formal charge
(Charge calculated from an atom in the Lewis
structure)
Useful to decide which Lewis structure is preferred when more
that one Lewis structure are possible. The less the formal charge
the more stable the molecule.
Formal Charge=(no. of Valence e- in the isolated atom) – (no. of bonds) – (no. of unshared e-)
Note that:
The sum of the formal charges of the atoms in a molecule or ion
must equal the charge on the molecule or ion.

22. Using formal charge to select Lewis structure
1) Construct Lewis structure that obey the octet rule.
2) Calculate the formal charge of this molecule
a)If the formal charges are zero, no need to look for another structure
b)If the formal charges are not zero, find another structure that do
not follow the octet rule by moving unshared pair of electrons
c) For Lewis structures having similar distributions of formal charges,
the most stable structure is the one in which negative formal charges
are placed on the more electronegative atoms.

23. Q) Which is the most likely Lewis structure for CH2O?

24. -1 +1
C = 4 ve-
O = 6 ve-
2H – 2x1 ve-
12 ve-
2 single bonds (2x2) = 4
1 double bond = 4
2 lone pairs (2x2) = 4
Total = 12 H C O
Formal Charge=(no. of Valence e- in the isolated atom) – (no. of bonds) – (no. of unshared e-)
formal charge on C
= 4 -
3 -
2 = -1
formal charge on O
= 6 -
3 -
2 = +1
= 1 -
1 -
0 = 0
formal charge on H

25. H C O
C = 4 ve-
O = 6 ve-
2H = 2x1 ve-
12 ve-
2 single bonds (2x2) = 4
1 double bond = 4
2 lone pairs (2x2) = 4
Total = 12
Formal Charge=(no. of Valence e- in the isolated atom) – (no. of bonds) – (no. of unshared e-)
formal charge on C
= 4 -
4 -
0 = 0
formal charge on O
= 6 -
2 -
4 = 0

26. the most likely Lewis structure for CH2O is :-1 +1 H C O

27. Q) Calculate the formal charge for the atoms in:1) 2) 3) 4)

28. Sa. Ex. : Write the possible Lewis structure for XeO3 and
show the most appropriate structure to the formal charges