Chemistry 101 : Chap. 8
Basic Concepts of Chemical Bonding
(1) 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
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)
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
Gilbert N. Lewis (1875-1946)
 This works only for representative
elements (main group)
Lewis Symbols
Elements
Group
e- Configuration
Lewis Symbol
Hydrogen
1A
1s1
H
Helium
8A
1s2
He
Lithium
1A
[He]2s1
Li
Berylium
2A
[He]2s2
Be
Boron
3A
[He]2s22p1
B
Carbon
4A
[He]2s22p2
C
Lewis Symbols
Elements
Group
e- Configuration
Nitrogen
5A
[He]2s22p3
N
Oxygen
6A
[He]2s22p4
O
Fluorine
7A
[He]2s22p5
F
Neon
8A
[He]2s22p6
Ne
Lewis Symbol
Note : All four sides of the symbol are equivalent
O
=
O
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
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
electron configuration:
K+ +
Cl
[Ar]
[Ar]
Both ions have an octet of electrons !
Ionic bonding
 Ionic Bonding : Cations (metals) and anions (non-metal)
combine to form ionic bonds
NaCl
Alternating positive and negative
charges
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?
Ionic bonding
The main driving force to form ionic bonds is the electrostatic
interaction between positive and negative ions.
Eel 
 Q Q
d
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
Ionic bonding
The stronger the ionic bond the higher the melting point
SrF2
+2, -2
r1
66, 133
1261oC
66, 140
2852oC
113, 133
1473oC
r2
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
F
+
F
H
H
F F
Each hydrogen has the
electron configuration of He
Each fluorine has the
electron configuration of Ne
Covalent bonding
 Lewis dot structure for covalent bonds
H
+
H
H
H
H
H
single covalent
bond
F
+
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)
Covalent bonding
 Example : Draw the Lewis dot structures of H2O and NH3
Covalent bonding
 Multiple bond
F
+
F
F F
or
F
F
Single bond
O
+
C
+
O
O
C
O
or
O
C
O
Double bond
N
+
N
N
N
or
N
N
Triple bond
Covalent bonding
 Single and Multiple bond
X
Distance between
atoms (bond
length) decreases
X
X
X
X
X
Bond strength
increases
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
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.
Drawing Lewis Structure
Lewis Structure of NH3
(1) Total number of valence electrons = 5 + 3  1 = 8
(2) Connect atoms with a single bond
and count the number electrons used
for single bond = 6
H
N
H
H
(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
Drawing Lewis Structure
Lewis Structure of CO
(1) Total number of valence electrons = 4 + 6 =10
(2) Connect atoms with a single bond
and count the number electrons used
for single bond = 2
C
(3) Complete the octets on the atoms bonded
to the central atom (6 electrons are used)
(4) Place remaining electrons
(10-2-6 = 2) on the central atom
(5) Carbon is NOT satisfying octet rule.
Need to have multiple bonds
C
C
C
O
O
O
O
C
O
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
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
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)?
Drawing Lewis Structure
Lewis Structure of ClO3- [ion]
(1) Total number of valence electrons = 7 + 63 + 1 = 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
(26-18-6 = 2) on the central atom
O
Cl
O
O
O
(5) All atoms are satisfying octet.
No need to consider multiple bonds
Cl
O
O
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
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 = 2 + 2  7 = 16
(2) Connect atoms with a single bond
and count the number electrons used
for single bond = 4
(3) Complete the octets on the atoms bonded
to the central atom (12 electrons are used)
Cl
Be
Cl
Cl
Be
Cl
(4) Place remaining electrons (16 - 4 -12 = 0) on the central atom : None left
(5) Be is not satisfying the octet rule, but no electron is available:
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 = 6 + 4  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
(3) Complete the octets on the atoms bonded
to the central atom (24 electrons are used)
S
F
(4) Place remaining electrons
(34-8-24 = 2) on the central atom
(5) S is not satisfying the octet rule
(10 electrons)
F
F
S
F
F
F
F
Drawing Lewis Structure :
Exceptions
 Molecule having an odd number of valence electrons :
Example = NO2
(1) Total number of valence electrons = 5 + 6  2 = 17
(2) Connect atoms with a single bond
and count the number electrons used
for single bond = 4
O
(3) Complete the octets on the atoms bonded
to the central atom (12 electrons are used)
(4) Place remaining electrons
(17-4-12 = 1) on the central atom
(5) Nitrogen has only 5 electrons.
Need to have multiple bonds
N
O
O
O
N
N
O
O
N
O
O
Free radical
Drawing Lewis Structure :
Exceptions
 Example : Determine the Lewis structure of BF3, BrF5 and OH
Drawing Lewis Structure :
Resonance
Lewis Structure of SO3
(1) Total number of valence electrons = 6 + 3  6 = 24
(2) Connect atoms with a single bond
and count the number electrons used
for single bond = 6
O
S
O
(3) Complete the octets on the atoms bonded
to the central atom (18 electrons are used)
O
S
(4) Place remaining electrons on the central atom.
No more electron is left (24-6-18=0)
O
(5) Sulfur is NOT satisfying octet rule.
S
Need to have multiple bonds
O
O
All three S-O bonds have the same length
O
O
O
O
O
O
S
S
O
O
Resonance structures
O
Drawing Lewis Structure :
Resonance
 Example : Determine the Lewis structure of O3 and HCO2-
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.
Properties of Covalent Bond
 Example : Predict which member of each set would have the
shortest bond length
SS
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.
H
H
C
H
H
H = 1660 kJ/mol
C (g)
+ 4 H (g)
C (g)
+ H (g) D (C-H) = 1660/4 kJ/mol
= 415 kJ/mol
(g)
per C-H bond:
C H (g)
Properties of Covalent Bond
Bond enthalpy can be used to estimate the enthalpy change
of chemical reactions, Hrxn
H2(g) + Cl2(g)  2HCl(g)
H1
H = ?
H2
Hrxn
Properties of Covalent Bond
H1 = D(H-H) + D(Cl-Cl) = 436kJ/mol + 243kJ/mol = 679 kJ/mol
H2 = 2 [  D(H-Cl) ] = 2  - 431 kJ/mol = - 862 kJ/mol
Hrxn = H1 + H2 = 697kJ/mol – 862 kJ/mol = -183 kJ/mol
Horxn = Σ n x Dbroken – Σ m x Dformed
moles of bonds
Properties of Covalent Bond
+
bonds broken
H–H
Cl – Cl
1
1
Bond Enthalpy (kJ/mol)
H–H
Cl – Cl
H – Cl
+
Hrxn
436
243
431
bonds formed
H – Cl
2
Hrxn = 1  436 + 1  243 – 2  431
= - 183 kJ/mol
Properties of Covalent Bond
 Example
CH4 (g)
: Estimate the Hrxn of following reaction
+
2 O2 (g)
→
CO2 (g)
+
2 H2O (g)
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)
Linus Carl Pauling
(1901-1994)
 most easily giving up electron
Electronegativity
Pauling scale of
electronegativity
Element
EN
F
4.0
O
3.5
Cl
3.0
N
3.0
C
2.5
H
2.1
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
Bond Polarity
 Polar covalent bond
When two different elements are bonded together,
there is unequal sharing of the electrons in the bond
+

H
Cl
polar covalent bond:
unequal sharing of electrons
The bonding pair of electrons
is pulled toward the chlorine
atom (partial charge)
Bond Polarity and Electronegativity
+

H
Na+
Cl
Cl -
polar covalent bond:
unequal sharing of electrons
ionic bond:
electrons are not shared
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
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