1. Chemistry 101 : Chap. 8 Basic Concepts of Chemical Bonding
Chemical Bonds, Lewis Symbols and the Octet Rule
(2) Ionic Bonding
(3) Covalent Bonding
(4) Bond Polarity and Electronegativity
(5) Drawing Lewis Structures
(6) Resonance Structures
(7) Exceptions to the Octet Rule
(8) Strengths of Covalent Bonds

2. Chemical Bonds Chemical bond is formed when two atoms or ions are held
together by the attractive force between them.
 Ionic Bond : a chemical bond formed between cation and anion
 Covalent Bond : a chemical bond formed between two nonmetallic
atoms by sharing one or more pairs of electrons.
 Metallic Bond : a chemical bond formed when valence electrons
of metal atom are attracted by the nuclei of surrounding atoms
(electrons are free to move throughout the metal)

3. Lewis Symbols
 Lewis electron dot structure (or Lewis symbol) : Symbol of
element surrounded by dots representing the valence electrons
in the atom
Lewis symbol for sulfur : [Ne]3s23p4 S
Maximum 2 electrons
on each side
 This works only for representative
elements (main group)
Gilbert N. Lewis ( )

4. Lewis Symbols Elements Group e- Configuration Lewis Symbol
Hydrogen A s H
Helium A s He
Lithium A [He]2s Li
Berylium A [He]2s Be
Boron A [He]2s22p B
Carbon A [He]2s22p C

5. Lewis Symbols O O = Elements Group e- Configuration Lewis Symbol
Nitrogen A [He]2s22p N
Oxygen A [He]2s22p O
Fluorine A [He]2s22p F
Neon A [He]2s22p Ne
Note : All four sides of the symbol are equivalent O = O

6. Lewis Symbols
 Elements in the same group of periodic table have the same
Lewis symbols
F Cl Br I
Elements in the same group have the same valence electron
configurations
For halogen atoms : ns2np5

7. Octet Rule
 Only the valence electrons are involved in chemical bonding.
 Octet Rule : When forming chemical bond, atoms tend to gain,
loose or share electrons in order to achieve a complete octet of
valence electrons (ns2np6)
 same electron configuration as noble gas atom + K + Cl K+ Cl
electron configuration:
[Ar]
[Ar]
Both ions have an octet of electrons !

8. Ionic bonding Ionic Bonding : Cations (metals) and anions (non-metal)
combine to form ionic bonds
NaCl
Alternating positive and negative
charges

9. Ionic bonding
 NaCl formation : Na(s) + ½ Cl2(g)  NaCl (s) Hof = -490 kJ
Metal : small ionization energy
Na(g)  Na+(g) + e IE = 496 kJ
Non-metal : large electron affinity
Cl(g) + e-  Cl-(g) EA = -349 kJ
Removing an electron from Na and transferring it to Cl
is NOT exothermic !
Then, why NaCl formation is an exothermic process?

10. Ionic bonding
The main driving force to form ionic bonds is the electrostatic
interaction between positive and negative ions.
charges of ions
distance between ions
 Strength of ionic bond depends on Eel
 the larger Eel, the stronger the bond
 the greater the charges, the stronger the bond
 the smaller the distance between the charges, the
stronger the bond

11. Ionic bonding The stronger the ionic bond the the melting point higher
66, 133
1261oC
66, 140
2852oC
SrF2
+2, -2
113, 133
1473oC r1 r2

12. Covalent bonding
 Covalent bond is formed when two atoms share electrons in order
to achieve the electron configuration of the nearest noble gas.
 satisfy octet rule H + H
Each hydrogen has the electron configuration of He H H
Each fluorine has the electron configuration of Ne F + F F

13. Covalent bonding  Lewis dot structure for covalent bonds H + H H H
single covalent
bond F + F F F F
A shared electron pair is drawn as a dash (two bonding electrons)
Unshared electrons are drawn as dots (lone-pair electrons)

14. Covalent bonding
 Example : Draw the Lewis dot structures of H2O and NH3

15. Covalent bonding  Multiple bond F + F F F or F F Single bond O + C O
Double bond N + N N or
Triple bond

16. Covalent bonding  Single and Multiple bond X X X X X X
Distance between atoms (bond length) decreases
Bond strength increases
X X
X X

17. Drawing Lewis Structure
 Things to know before you start to draw Lewis structure
 Chemical formulas are often written in the order in which the atoms
are connected
ex) HCN
 Hydrogen has only two electrons (shared) and always has only
one covalent bond
 The central atom is usually written first
ex) NH3, CCl4, CHCl3, PCl3

18. Drawing Lewis Structure
 Rules for drawing Lewis structure
(1) sum the number of valence electrons from all atoms
(2) write the symbols for the atoms and connect them with a single bond
(3) complete the "octet rule" for the atoms bonded to central atom
(4) place any left over electrons on the central atom
(5) If there are not enough electrons to give the central atom
8 electrons, try multiple bonds.

19. Drawing Lewis Structure
Lewis Structure of NH3
(1) Total number of valence electrons =  1 = 8
(2) Connect atoms with a single bond
H N H H
and count the number electrons used
for single bond = 6
(3) Complete the octets on the atoms bonded to the central atom : done
(4) Place remaining electrons
(8-6=2) on the central atom
H N H H
(5) All atoms are satisfying octet. No need to consider multiple bonds

20. Drawing Lewis Structure
Lewis Structure of CO
(1) Total number of valence electrons = =10
(2) Connect atoms with a single bond
C O
and count the number electrons used
for single bond = 2
(3) Complete the octets on the atoms bonded
to the central atom (6 electrons are used)
C O
(4) Place remaining electrons
( = 2) on the central atom
C O
(5) Carbon is NOT satisfying octet rule.
Need to have multiple bonds
C O
C O

21. Drawing Lewis Structure
 Example : Determine the Lewis structure of HCN
(1) Total number of valence electrons
(2) Connect atoms with a single bond
and count the number electrons used
for single bond
(3) Complete the octets on the atoms bonded
to the central atom
(4) Place remaining electrons on the central atom.
(5) Carbon is NOT satisfying octet rule.
Need to have multiple bonds

22. Drawing Lewis Structure
 Example : Determined the Lewis structure of CH2O
(1) Total number of valence electrons
(2) Connect atoms with a single bond
and count the number electrons used
for single bond
(3) Complete the octets on the atoms bonded
to the central atom
(4) Place remaining electrons on the central atom.
(5) Carbon is NOT satisfying octet rule.
Need to have multiple bonds

23. Drawing Lewis Structure
 Example : Determined the Lewis structure of H2O2
(1) Total number of valence electrons
(2) Connect atoms with a single bond
and count the number electrons used
for single bond
(3) Complete the octets on the atoms bonded
to the central atom
(4) Place remaining electrons on the central atom.
(5) All atoms are satisfying octet.
No need to consider multiple bonds
What happens if you choose a different geometry in step (2)?

24. Drawing Lewis Structure
Lewis Structure of ClO3- [ion]
(1) Total number of valence electrons = 7 + 6 = 26
(2) Connect atoms with a single bond
and count the number electrons used
for single bond = 6
O Cl O O
(3) Complete the octets on the atoms bonded
to the central atom (18 electrons are used)
O Cl O O
(4) Place remaining electrons
( = 2) on the central atom
O Cl O O
O Cl O O
(5) All atoms are satisfying octet.
No need to consider multiple bonds

25. Drawing Lewis Structure
 Example : Determined the Lewis structure of ClO2-
(1) Total number of valence electrons
(2) Connect atoms with a single bond
and count the number electrons used
for single bond
(3) Complete the octets on the atoms bonded
to the central atom
(4) Place remaining electrons
on the central atom
(5) All atoms are satisfying octet.
No need to consider multiple bonds

26. Drawing Lewis Structure : Exceptions
Atoms having fewer than 8 valence electrons :
Group IIA and IIIA (mostly Be, B).
Example = BeCl2
(1) Total number of valence electrons =  7 = 16
(2) Connect atoms with a single bond
and count the number electrons used
for single bond = 4
Cl Be Cl
(3) Complete the octets on the atoms bonded
to the central atom (12 electrons are used)
Cl Be Cl
(4) Place remaining electrons ( = 0) on the central atom : None left
(5) Be is not satisfying the octet rule, but no electron is available:

27. Drawing Lewis Structure : Exceptions
Atoms having more than 8 valence electrons :
central atom with n 3, which can use d-orbitals for bonding
Example = SF4
(1) Total number of valence electrons =  7 = 34
(2) Connect atoms with a single bond
and count the number electrons used
for single bond = 8
F F S
F F
F F
(3) Complete the octets on the atoms bonded
to the central atom (24 electrons are used) S
F F
(4) Place remaining electrons
( = 2) on the central atom
F F
(5) S is not satisfying the octet rule
(10 electrons) S
F F

28. Drawing Lewis Structure : Exceptions
Molecule having an odd number of valence electrons :
Example = NO2
(1) Total number of valence electrons =  2 = 17
(2) Connect atoms with a single bond
and count the number electrons used
for single bond = 4
O N O
(3) Complete the octets on the atoms bonded
to the central atom (12 electrons are used)
O N O
(4) Place remaining electrons
( = 1) on the central atom
O N O
(5) Nitrogen has only 5 electrons.
Need to have multiple bonds
O N O
Free radical

29. Drawing Lewis Structure : Exceptions
 Example : Determine the Lewis structure of BF3, BrF5 and OH

30. Drawing Lewis Structure : Resonance
Lewis Structure of SO3
(1) Total number of valence electrons =  6 = 24
(2) Connect atoms with a single bond
and count the number electrons used
for single bond = 6 O S O O O
(3) Complete the octets on the atoms bonded
to the central atom (18 electrons are used) S O O
(4) Place remaining electrons on the central atom.
No more electron is left ( =0) O O O
(5) Sulfur is NOT satisfying octet rule.
Need to have multiple bonds S S S O O O O O O
All three S-O bonds have the same length
Resonance structures

31. Drawing Lewis Structure : Resonance
 Example : Determine the Lewis structure of O3 and HCO2-

32. Properties of Covalent Bond
 Bond length : The distance between two bonded atoms
 
bond length
Bond length depends on the size of two atoms and the
number of covalent bond (single, double or triple) between them.

33. Properties of Covalent Bond
 Example : Predict which member of each set would have the
shortest bond length S S

34. Properties of Covalent Bond
Bond Enthalpy: Energy required to completely separate two
bonded atoms in gas phase. A short bond is usually harder to break.
C (g) H (g)
DH = 1660 kJ/mol
(g)
per C-H bond:
C H (g)
C (g) H (g)
D (C-H) = 1660/4 kJ/mol
= 415 kJ/mol

35. Properties of Covalent Bond
Bond enthalpy can be used to estimate the enthalpy change
of chemical reactions, Hrxn
H2(g) + Cl2(g)  2HCl(g) H = ?
H1
H2
Hrxn

36. Properties of Covalent Bond
H1 = D(H-H) + D(Cl-Cl) = 436kJ/mol + 243kJ/mol = 679 kJ/mol
H2 = 2 [  D(H-Cl) ] = 2  kJ/mol = kJ/mol
Hrxn = H1 + H2 = 697kJ/mol – 862 kJ/mol = -183 kJ/mol
DHorxn = Σ n x Dbroken – Σ m x Dformed
moles of bonds

37. Properties of Covalent Bond+ +
Hrxn
bonds broken
H – H
Cl – Cl
bonds formed
H – Cl
Bond Enthalpy (kJ/mol)
H – H
Cl – Cl
H – Cl
Hrxn = 1   243 – 2  431
= kJ/mol

38. Properties of Covalent Bond
 Example : Estimate the Hrxn of following reaction
CH4 (g) O2 (g) → CO2 (g) H2O (g)

39. Electronegativity
 Electronegativity : A measure of the attraction an atom has
for the electron in a bond
Metals  low electronegativity
Nonmetals  high electronegativity
electronegativity scale:
Fluorine = 4 (most electronegative)
 most strongly attracting electron
Cesium = 0.7 (least electronegative)
 most easily giving up electron
Linus Carl Pauling
( )

40. Electronegativity Pauling scale of electronegativity Element EN F 4.0Cl N C H

41. non-polar covalent bond:
Bond Polarity
 Nonpolar covalent bond
When two atoms of same element are bonded together,
there is equal sharing of the electrons in the bond Cl Cl
non-polar covalent bond:
equal sharing of electrons

42. unequal sharing of electrons
Bond Polarity
 Polar covalent bond
When two different elements are bonded together,
there is unequal sharing of the electrons in the bond +
 -
The bonding pair of electrons
is pulled toward the chlorine
atom (partial charge) H Cl
polar covalent bond:
unequal sharing of electrons

43. Bond Polarity and Electronegativity+
 -
Na+ Cl - H Cl
ionic bond:
electrons are not shared
polar covalent bond:
unequal sharing of electrons
EN = 3.0 – 2.1 = 0.9
EN = 3.0 – 0.9 = 2.1
EN < 0.5
non-polar bond
0.5  EN < 2.0
polar bond
EN  2.0
ionic bond

44. Bond Polarity and Electronegativity
 Example : For each pair of bonds, predict which bond is more polar
and the partial charge on the atoms
(a) Cl – Br Br – F
(c) C – H C – O
(b) O – F S – F
(d) H – O Na – O