This document discusses covalent bonds and Lewis structures. It begins by listing learning objectives related to covalent bonding concepts. It then defines key terms like Lewis structure, covalent bond, and electronegativity. The document goes on to explain the formation of covalent bonds via electron sharing. It discusses exceptions to the octet rule. Guidelines are provided for drawing Lewis structures, including dealing with resonance. Various examples of Lewis structures are worked through. The document concludes with a review of naming covalent compounds.

1. COVALENT BONDS &
LEWIS STRUCTURE
Prepared by:
Mrs. Eden C. Sanchez

2. Learning Objectives
1. Illustrate the formation of covalent bonds in
terms of electron sharing
2. Apply the octet rule in forming covalent
compounds
3. Define electronegativity
4. Describe the electronegativity trends in the
periodic table
2

3. Learning Objectives
5. Draw Lewis structure of covalent
compounds
6. Identify lone pairs and bond pairs;
7. Draw the resonance structures of covalent
compounds
8. Determine the polarity of a bond based on
the electronegativities of the bonding
atoms
3

4. Learning Objectives
9. Determine whether a bond is ionic, polar
covalent, or covalent based on the
differences in electronegativities of the
bonding atoms
4

5. Keywords
a. Lewis structure
b. Covalent bond
c. Lone pair
d. Bond pair
e. Single bond
f. Double bond
5

6. Keywords
g. Triple bond
h. Nonpolar covalent bond
i. Polar covalent bond
j. Electronegativity
k. Percent ionic character
l. Resonance
6

7. Keywords
m. Incomplete octet
n. n. Expanded octet
7

8. FORMATION OF THE
COVALENT BOND
▰ Gilbert Lewis suggested that the chemical
bond is formed by sharing of electrons in
atoms.
▰ Example:
▰ The two electrons are shared equally
between the two atoms forming a covalent
bond.
8

9. FORMATION OF THE
COVALENT BOND
▰ The electrons are attracted to the nuclei of
both atoms keeping the atoms together to
form a molecule.
▰ formation of the covalent bond for the F2
molecule
9

10. FORMATION OF THE
COVALENT BOND
▰ The representation of the covalent
compound is called the Lewis structure. In
the Lewis structure, shared electrons that
form a bond is represented by a line or a
pair of dots; lone pairs are represented by
dots above the atom.
10

11. FORMATION OF THE
COVALENT BOND
a. From the Lewis structure of F2, how many
electrons are around each fluorine atom in
F2?
b. How many bond pairs are there in the F2
molecule?
11

12. FORMATION OF THE
COVALENT BOND
c. How many lone pairs are there in the F2
molecule?
d. Illustrate the formation of the covalent
bond in Cl2. How many bond pairs are
there?
e. How many lone pairs?
12

13. FORMATION OF THE
COVALENT BOND
c. Illustrate the formation of the covalent
bond in HCl.
13

14. Exercises:
1. Draw the Lewis structure for H2O, CH4
(methane), and for NH3.
2. Which of the three molecules has the
largest number of bond pairs (covalent
bonds)?
3. Draw the Lewis structure for carbon
dioxide, CO2.
14

15. ▰ The examples of CO2 and N2 show that
there are different types of covalent bonds
that are formed.
▰ Single bonds are formed when two atoms
are held together by one pair of electrons.
▰ Multiple bonds can be formed also
15

16. ▰ A double bond is from the sharing of two
pairs of electrons such as in the case of O
and C in CO2.
▰ A triple bond exists in N2 where the two N
atoms are held by three pairs of electrons.
16

17. Types of Covalent Bond
▰ Experimental evidence has shown that
electrons are not equally shared between H
and F; the electrons spend more time near F
rather than H. Therefore the electron density
is shifted more towards F rather than H.
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18. ▰ This leaves the F end of the molecule
partially negative, δ - , and the H end of
the molecule partially positive, δ+, such
bond is referred to as a polar covalent
bond.
18

19. ▰ The polar covalent bond is somewhere
between a purely covalent (nonpolar) bond
and an ionic bond (where there is almost
complete transfer of electrons).
19

20. Electronegativity
▰ property that distinguishes the polarity of
bonds
▰ the tendency of an atom in a chemical bond
to attract electrons toward itself
▰ is a theoretical concept and devised as a
relative scale. That is, it can be estimated
relative to, or in comparison to, other
elements in chemical bonds.
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21. Electronegativity
▰ Linus Pauling developed a relative scale of
electronegativities which is widely used in
General Chemistry
▰ In contrast, ionization energies and electron
affinities are physically measurable
properties of elements.
21

22. 22

23. ▰ The difference in the electronegativity values
( EN) of two bonded atoms determines the
percent ionic character of the bond. If the
bond is between two identical elements, like
F—F, then the bond is purely covalent with
0 percent ionic character. The difference in
electronegativity is 0.
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24. ▰ For the molecule H—Cl, the difference in
electronegativity is 0.9 showing that the bond
is a polar covalent bond. A 50% ionic
character corresponds to EN=1.7. While
there is no bond that is 100% ionic, an
electronegativity difference of 2.0 or greater
is usually classified to be predominantly ionic.
24

25. ▰ When EN ≥ 2.0, the bond is
predominantly ionic.
25

26. Exercises:
I. Classify the following bonds as ionic, polar
covalent, or covalent.
1. The C-C bond in H3CCH3
2. The K-I bond in KI
3. The C-F bond in CF4
4. The N-H bond in NH3
26

27. Exercises:
II. Arrange the following bonds according to
increasing bond polarity: Cs to F, Cl to Cl, Br
to Cl, Si to C.
27

28. Guidelines in Writing the Lewis
Structure of Covalent Molecules
1. Draw a skeletal structure of the molecule
putting bonded atoms next to each other.
In general, the least electronegative atom
occupies the central position. H and F
usually occupy terminal (end) positions.
28

29. Guidelines in Writing the Lewis
Structure of Covalent Molecules
2. Count the total number of valence
electrons from all the atoms in the
structure. Add electrons corresponding to
the charge for negative ions; subtract
electrons corresponding to the charge for
positive ions.
29

30. Guidelines in Writing the Lewis
Structure of Covalent Molecules
3. Distribute the valence electrons to the non-
central atoms such that these atoms fulfill
the octet rule. Remaining electrons are
assigned to the central atom. Remember
that bonds are equivalent to 2 electrons.
4. If the valence electrons are not enough,
multiple bonds may be formed.
30

31. Exercises
1. Write the Lewis structure for NCl3.
2. Write the Lewis structure of OCS. C is the
central atom.
3. Write the Lewis structure of CN–.
31

32. Exercises
4. Write the Lewis structure of the following
molecules:
a. Ethylene, C2H4
b. Acetylene, C2H2
c. Carbon tetrachloride, CCl4
d. COBr2 (for the skeletal structure, C is
bonded to O and Br atoms)
32

33. Lewis Structure & Resonance
▰ Write the Lewis structure for the ozone
molecule, O3.
▰ To resolve this discrepancy, we represent
the ozone molecule using the two structures
presented as:
33

34. Lewis Structure & Resonance
▰ The above structures is called a resonance
structure. The double sided arrow shows
that the structures are resonance
structures. A resonance structure is one of
two or more Lewis structures for a molecule
that cannot be represented accurately by
only one Lewis structure.
34

35. Exercise:
▰ Draw the resonance structures for the
carbonate ion, CO3
2- .
35

36. EXCEPTIONS TO THE OCTET
RULE:
The octet rule works best for second-period
elements. Hence there are many exceptions.
They fall into three categories:
a. Incomplete octet
b. Odd number of electrons
c. Expanded Octet
36

37. EXCEPTIONS TO THE OCTET
RULE:
▰ Incomplete octet
An example of a molecule with incomplete
octet is BeH2, beryllium hydride.
H – Be – H
▰ There are only 4 electrons around Be and
not 8. Boron and aluminum also form
molecules with incomplete octets.
37

38. EXCEPTIONS TO THE OCTET
RULE:
▰ Draw the Lewis structure of aluminum
triiodide, AlI3, showing the incomplete octet.
38

39. EXCEPTIONS TO THE OCTET
RULE:
▰ Molecules with Odd Number of
Electrons
▰ Examples are nitric oxide, NO, and
dinitrogen dioxide, N2O.
39

40. EXCEPTIONS TO THE OCTET
RULE:
▰ The odd numbered molecules are
sometimes referred to as radicals. They
are generally highly reactive.
40

41. EXCEPTIONS TO THE OCTET
RULE:
▰ Expanded Octets
Atoms belonging to the second period
cannot have more than eight valence
electrons around the central atom because
they only have the 2s and 2p subshells. This
is different for atoms of elements in the 3rd
period and beyond.
41

42. EXCEPTIONS TO THE OCTET
RULE:
▰ These elements have 3d orbitals that can
participate in the bonding. Hence they can
have more than eight valence electrons
around the central atom. An example is
SF6, sulfur hexafluoride and phosphorus
pentafluoride, PF5.
42

43. EXCEPTIONS TO THE OCTET
RULE:
43

44. EXCEPTIONS TO THE OCTET
RULE:
▰ Another example is BrF5
44

45. NAMING COVALENT
COMPOUNDS: (A REVIEW)
1. For binary compounds, state the name of
the first element. The name of the second
element ends in –ide.
HF Hydrogen fluoride
HI Hydrogen iodide
SiC Silicon carbide
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46. NAMING COVALENT
COMPOUNDS: (A REVIEW)
2. Prefixes are used to denote the number of
atoms in the formula. The prefix “mono”
usually omitted for the first element in the
formula.
CO carbon monoxide
CO2 carbon dioxide
NO2 nitrogen dioxide
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47. NAMING COVALENT
COMPOUNDS: (A REVIEW)
N2O4 dinitrogen tetroxide
CCl4 carbon tetrachloride
SF6 sulfur hexafluoride
47

48. Seatwork
▰ Look for at least 2 examples of covalent
compounds that can be found in nature or
used in everyday life. Include the following
information:
a. uses of the covalent compound
b. chemical formula and chemical name of
the covalent compound
c. structure of the compound
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49. 49
Thank You!
Any questions?

50. CREDITS
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