Chemical Bonding
Chapter 6
Sections 1, 2, and 5
Chemical Bonds
A chemical bond is the mutual electrical attraction
between the nuclei and valence electrons of
different atoms that bind the atoms together
Noble gases tend not to do this because of their filled
s and p orbitals.
They have a stable octet: outer s and p orbitals are
completely filled with e-’s ( 8 total)
Chemical Bonds
Atoms that don’t have a stable octet are more
reactive because their potential energy is higher.
They become more stable by decreasing their
potential energy.
Octet Rule: chemical compounds tend to form so that
each atom has an octet of e-’s in its highest
occupied energy level
How to do this? Gain , lose or share electrons between
atoms
Chemical Bonds
By forming a chemical bonds, atoms gain stability!
Chemical changes always involve energy
What type of bonds can be formed?
Ionic bond
Covalent bond
Nonpolar covalent
Polar covalent
Chemical Bonds
Ionic bonding: bonds that result from electrical
attractions between cations and anions
Covalent bonding: sharing of electron pairs between 2 or
more atoms
*** In reality, bonding is often somewhere between the
two extremes***
Two types of Covalent Bonds
Nonpolar –covalent: equal sharing of electron pairs
Polar-covalent: unequal attraction for the shared
electrons
How can we determine the type of bond?
Knowing how strong an atom’s ability is to attract
electrons (aka electronegativity), helps us
determine if it will form a ionic or covalent bond
with another atom.
A large difference in E.N. between atom’s will result
in an ionic bond
A small difference between atom’s will result in a
form of covalent bonding
What type of Bond is it?
Electronegativity
Difference
Bond Type
0 to 0.3
Nonpolar Covalent
0.4 to 1.7
Polar Covalent
 1.7
Ionic
Increasing difference in electronegativity
Nonpolar
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
Do you see any trends?
A metal and nonmetal tend to form ionic compounds
Nonmetal and nonmetal tend to form polar-covalent or
nonpolar- covalent compounds
Polar covalent bond or polar bond :
covalent bond with greater electron density
around one of the two atoms
electron poor
region
H
electron rich
region
F
e- poor e- rich
H
F
d+ d-
Classify the following bonds as ionic, polar
covalent,or covalent:
CsCl
Cs – 0.7
Cl – 3.0
3.0 – 0.7 = 2.3
Ionic
H2S
H – 2.1
S – 2.5
2.5 – 2.1 = 0.4
Polar Covalent
N – 3.0
3.0 – 3.0 = 0
Nonpolar
Covalent
N2
N – 3.0
Properties of Molecular
Covalent Compounds
Not very soluble in water
Do not conduct electricity
Low melting points
Low boiling points
Can be solids, liquids and gases at room temperature
Comparison of Ionic and Covalent Compounds
Types of Crystals
Lewis Structures
Review:
What are valence electrons?
Lewis Dot Diagrams
− an electron-configuration notation with only the valence
electrons of an element are shown, indicated by dots
placed around the element’s symbol.
− the inner core electrons are not shown.
The Octet Rule
•Eight electrons in the valence shell (filling s and p
orbitals) make an atom STABLE
2
6
sp
This is called the octet rule
•Bond formation follows the octet rule: Chemical
compounds tend to form so that each atom:
by gaining, losing, or sharing electrons, has an octet
of electrons in its valence energy level.
Lewis Structures for Compounds
•The pair of dots between two symbols
represents the shared pair of a covalent bond.
•Each fluorine atom is surrounded by three pairs
of electrons that are not shared in bonds.
•An unshared pair, also called a lone pair, is a pair
of electrons that is not involved in bonding and
that belongs exclusively to one atom.
F F
Lewis Structures
•The pair of dots representing a shared pair of
electrons in a covalent bond is often replaced by
a long dash.
H
H
F
F
covalent bond : is a chemical bond in which
two or more electrons are shared by two
atoms.
Why should two atoms share electrons?
F
+
7e-
F
F F
7e-
8e- 8e-
Lewis structure of F2
single covalent bond
lone pairs
F
F
single covalent bond
lone pairs
F F
lone pairs
lone pairs
Multiple Covalent Bonds
•double covalent bond or double bond :
covalent bond in which two pairs of electrons
are shared between two atoms
•shown by two side-by-side pairs of dots or by
two parallel dashes
H
H
C C
H
H
H
or
H
C C
H
H
Multiple Covalent Bonds
•triple covalent bond or triple bond :
covalent bond in which three pairs of electrons
are shared between two atoms.
N
H C
N or N N
C H or H C
C H
Single Bond – two atoms share one pair of electrons
single covalent bonds
H
+
O +
H
H O H
or
H
O
H
2e-8e-2eDouble Bond – two atoms share two pairs of electrons
O C O
or
O
O
C
double bonds
- 8e8e- 8ebonds
double
Triple Bond – two atoms share three pairs of electrons
N N
triple
bond
8e-8e
or
N
N
triple bond
Lengths of Covalent Bonds
Bond
Type
Bond
Length
(pm)
C-C
154
CC
133
CC
120
C-N
143
CN
138
CN
116
Bond Lengths
Triple bond < Double Bond < Single Bond
Writing Lewis Structures
1. Draw skeletal structure of compound showing
what atoms are bonded to each other. Put least
electronegative element in the center.
2. Count total number of valence e-. Add 1 for
each negative charge. Subtract 1 for each
positive charge.
3. Complete an octet for all atoms except
hydrogen
4. If structure contains too many electrons, form
double and triple bonds on central atom as
needed.
Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete
octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
F
N
F
F
Write the Lewis structure of the carbonate ion (CO32-).
Step 1 – C is less electronegative than O, put C in center
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)
-2 charge – 2e4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete
octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-
O
C
O
O
2 single bonds (2x2) = 4
1 double bond = 4
8 lone pairs (8x2) = 16
Total = 24
resonance structure: one of two or more Lewis structures for
a single molecule can be drawn to represent a molecule
O
O
+
-
-
O
O
+
O
O
What are the resonance structures of the
carbonate (CO32-) ion?
-
O
C
O
O
-
O
C
O
O
-
-
-
O
C
O
O
-
Exceptions to the Octet Rule
The Incomplete Octet
BeH2
BF3
B – 3e3F – 3x7e24e-
Be – 2e2H – 2x1e4e-
F
B
F
H
F
Be
H
3 single bonds (3x2) = 6
9 lone pairs (9x2) = 18
Total = 24
Exceptions to the Octet Rule
Odd-Electron Molecules
NO
N – 5eO – 6e11e-
N
O
The Expanded Octet (central atom with principal quantum number n > 2)
SF6
S – 6e6F – 42e48e-
F
F
F
S
F
F
F
6 single bonds (6x2) = 12
18 lone pairs (18x2) = 36
Total = 48
Molecular Geometry
Chapter 6.5
VSEPR THEORY
Lewis Dot Diagrams are 2D but we live in a 3D world.
How are molecules actually arranged??
Follows the Valance Shell Electron Pair Repulsion Theory or
VSEPR
AB2 – Linear
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
2
0
180˚
Cl
Be
Cl
AB3 – Trigonal Planar
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
3
0
120˚
AB2E1 – Bent
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
2
1
<120˚
AB4 – Tetrahedral
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
4
0
109.5˚
AB3E1 – Trigonal
Pyramidal
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
3
1
107˚
AB2E2 – Bent
Number of Surround
Atoms
Number of Lone Pairs
Bond Angle
2
2
104.5˚
Predicting Molecular
Geometry
1. Draw Lewis structure for molecule.
2. Count number of lone pairs on the central atom and
number of atoms bonded to the central atom.
3. Use VSEPR to predict the geometry of the molecule.
What are the molecular geometries of SO2 and SF4?
O
S
AB2E
bent
F
O
F
S
F
AB4E
F
distorted
tetrahedron
The electronegativity of an atom will
create a dipole, or polar molecule.
electron poor
region
electron rich
region
H
F
d+
d-
Which of the following molecules have a dipole moment?
H2O, CO2, SO2, and CH4
O
S
dipole moment
polar molecule
dipole moment
polar molecule
H
O
C
O
no dipole moment
nonpolar molecule
H
C
H
H
no dipole moment
nonpolar molecule
Intermolecular forces: attractive forces
between molecules.
Intramolecular forces: hold atoms together,
attractive forces within a molecule.
Generally, intermolecular forces are much
weaker than intramolecular forces.
Properties of Ionic
Compounds
Combination of ions (cation/anion)
Hard and Brittle
Tightly packed solids in a crystal lattice
Usually soluble in water
Conducts electricity when dissolved
High melting points
Breaking Ionic Bonds
Ionic Bonds are very tightly bound
positive and negative attraction
A LOT of energy needs to be put in to break an ionic bond
How does this affect the melting point?