Chapter 8 – Covalent Bonding
8.1 The Covalent Bond
8.2 Naming Molecules
8.3 Molecular Structures
8.4 Molecular Shape
8.5 Electronegativity and Polarity
Topic is Lewis Structures (combination
of material found in 2 sections)
Sections 8.1/8.3 Covalent Bond/Molecular
Structures
Atoms gain stability when they share
electrons and form covalent bonds. The
sharing can be described by the Lewis
structure of the compound.
• Apply the octet rule to atoms that form covalent
bonds.
• Describe the similarities and differences between
ionic and covalent bonding.
• Describe the forces that act and energy changes that
occur when atoms form a molecule.
Sections 8.1/8.3 Covalent Bond/Molecular
Structures
• Categorize valance electrons as bonding or nonbonding.
• List the basic steps used to draw Lewis structures.
• Describe the formation of single, double, and triple
covalent bonds using Lewis structures.
• Explain why resonance occurs, and identify
resonance structures.
• Draw valid Lewis structures for molecules, including
those involving multiple bonds, resonance, expanded
octets, and electron deficient molecules.
Why Atoms Bond
Noble gas configuration especially
stable
• ns2np6 (except for He)
• Full outer energy level
• Octet of electrons
Atoms bond to achieve a lower energy
state (more stability)
Ionic vs Covalent Bonding
In ionic bonding, electrons transferred
to achieve octet for each ion
• Number of ionic compounds small
compared to total number of known
compounds
In covalent bonding, electrons shared
to achieve octet (mostly) for each atom
Covalent Bonding
In covalent bonding, electrons shared
to achieve octet (mostly) for each atom
• Sharing occurs when electronegativities
of atoms same or similar
• Majority of covalent bonds formed
between nonmetallic elements
• Electronegativity difference < 1.7 (see
next slide) – bond will have more
covalent character than ionic character
100
75
Ionic Bonds
50
25
% Ionic Character
EN Difference & Bond Character
0
Covalent
Bonds
1.0
2.0
3.0
Electronegativity Difference
Covalent Bonding
In covalent bonding, electrons shared
to achieve octet (mostly) for each atom
Molecule formed when 2 or more atoms
bond covalently
Covalent Bonding – Forces
Covalent Bonding – Forces
No
interaction
Nucleus
attracted
to other
atoms
electrons
–
Not
optimum
distance
Nucleus
attracted
to other
atoms
electrons
–
optimum
distance
Net
repulsion
from
positive
nuclei
Potential Energy (kJ/mol)
Covalent Bonding – Energy for H2
-432 kJ/mol
100
200
Internuclear Distance (pm)
Lewis Structures - Atoms
(Electron Dot Diagrams)
Way of keeping track of
valence electrons
To write for atom
• Write symbol for element
• Put one dot for each
valence electron
• Don’t pair up until you
have to (Hund’s rule)
X
Lewis Structure - Covalent Molecules
Valence electrons of each element in
molecule are divided into 2 categories:
• Bonding – pair of electrons shared by
two atoms to form the covalent bond

Shared pair represented by a line
connecting the element symbols H—H
• Nonbonding – called lone pairs

A few molecules have odd # total
electrons – have unpaired nonbonding
electron
Lewis Structure - Covalent Molecules
Example – formation of H2 molecule
H does not
form octet
Bond = shared electron pair
Space-Filling Model View
Formation of H2

+
H•
•H
H••H
_
Bond = shared electron pair H H
Ways of Representing Molecules: H2O
Orbital Model
Ball-and-Stick Model
Structural Formula
Space-Filling Model
Ways of Representing Molecules: PH3
Covalent Bonding – F2
F 1s22s22p5
7 valence electrons
Forms F2 molecule
Each F shares 1 valence electron
Molecule is more stable than individual
atoms
Lewis Structure - Covalent Molecules
Example – formation of F2 molecule
Bond = shared electron pair
Octet
formed
Lewis Structure - Covalent Molecules
Example – formation of F2 molecule
Octet
formed
Lewis Structure - Covalent Molecules
Example – formation of H2O molecule
Two lone
pairs
Shape of
Octet
molecule
formed
Bonds = shared electron pairs
Lewis Structure - Covalent Molecules
Example – formation of ammonia, NH3
Lone
pair
+
+
+

Bonds =
Shape
of
shared
molecule
electron
pairs
Octet
formed
Multiple Covalent Bonds
C, N, O, S often form multiple bonds
Double bond – O2 (6 valence e per O)
+

Triple bond – N2 (5 valence e per N)
+

Guide for Writing Lewis Structures
Similar to procedure on p. 254, but without # of bonding pairs
Step 1 – Write skeletal structure
• least electronegative atom usually
occupies central position
Step 2 – Count total number of valence
electrons
• polyatomic anions, add # of - charges

e.g. CO32- add 2 electrons to total
• polyatomic cations, subtract # of +
charges
Step 3 – Place single bond between
central atom and surrounding atoms
Step 4 – Complete octet for terminal
atoms (not for H)
Step 5 – Add remaining to central atom
Step 6 – If octet rule not satisfied for
central atom, add multiple (double,
triple) bonds between terminal and
central atom, using the lone pairs from
the terminal atoms
Lewis Structures – Common Bonding Patterns
C 4 bonds & 0 lone pairs
4 single (CH4), or 2 double (CO2), or single +
triple (HCCH), or 2 single + double (CH2CH2)
N 3 bonds & 1 lone pair (NH3)
O 2 bonds & 2 lone pairs (H2O)
H & halogen 1 bond (CH4, CF4)
Be 2 bonds & 0 lone pairs (BeH2, electron def.)
B 3 bonds & 0 lone pairs (BH3, electron def.)
B
C
N
O
F
Lewis Structure Examples
Total
Valence
Electrons
a, HF
Draw
Single
Bonds
1+7= 8
H-F
Calculate
Number of
Electrons
Remaining
Use
Remaining
Electrons to
Achieve
Noble Gas
Configuration
6
Check
Number of
Electrons
H, 2
H F
F, 8
b, N2
5 + 5 = 10
c, NH3
5 + 3(1) = 8
d, CH4
e, CF4
4 + 4(1) = 8
4 + 4(7) = 32
N-N
H
H
F
N H
8
2
H
H
H
H
C H
0
H
f,
5 + 6 - 1 = 10
F
F
N-O
H, 2
C, 8
H
C F
H, 2
N, 8
C H
H
24
N,8
N H
H
F, 8
F
C
F
C, 8
F
F
NO+
N
N
8
N
O
+
N, 8
O, 8
Practice
Problems 1-5 page 244
Problems 37-38, page 255
Problems 39-40, page 256 (mult bonds)
Problems 41-42 page 257 (ions)
Problems 104(a-d), page 275
Problems 1(a-d), page 979
Problems 4(a-e) page 980
Lewis Structure Example: NO3─
1. Write skeletal structure
N central because it is
least electronegative
O
O
2. Count valence electrons
N=5
3O = 3 x 6 = 18
(-) = 1
Total = 24 e-
N
O
Example NO3
3. Attach atoms with
single bonds (pairs
of electrons) &
subtract from total
Electrons
Start 24
Used 6
Left
18
─,
Continued
O

O —
N
— O
Example NO3─ , Continued
4. Complete octets,
outside-in
Keep going until all
atoms have an octet or
you run out of electrons
Electrons
Start 18
Used 18
Left
0
:
:

O :



O — N — O


:
Example NO3
─,
5. If central atom does
not have octet, bring in
electron pairs from
outside atoms to share
If structure is an ion,
use brackets and
indicate the charge
6. For this ion an extra
step is needed – draw
resonance structures
Continued
:
:

O :



O — N — O


:
:

O

|
O — N

:
-1
:

O:
Example NO3─ , Continued
Can have more
than one correct
Lewis structure
for molecules or
ions with double
and single bonds
-
Resonance Structures
Resonance structures
differ only in position of
electron pairs, never the
atom positions
Molecule behaves as if it
had only one structure
(the average one)
• NO3- has all bond lengths
identical
-
Practice (Resonance Structures)
Problems 43-46 page 258
Problems 101,103 pages 274-5
Problems 5, 6 page 980
Practice—Lewis Structures
NClO
H3PO4
H3BO3
SO3-2
NO2-1
P2H4
Practice—Lewis Structures
NClO
18 e-
*
••
N
••
•O
•
H3BO3
24
e- H
NO2-1
18
e-
••
O
••
••
•O
•
H3PO4
••
Cl ••
••
••
•O
•
B
*
••
N
H
••
O
••
••
O ••
••
32 eSO3-2
H
-1
26 e-
H
14 e-
••
O
••
••
•O
•
••
P2H4
H
* Has resonance structures
••
•O•
• •
P
•O
•
••
••
•O•
• •
S
••
H
H
P
••
P
••
••
O
••
H
H
-2
••
O ••
••
H
Exceptions to Octet Rule
Molecules with odd number of total
valence electrons
NO2 – 17 valence electrons
Also ClO2, NO
Exceptions to Octet Rule
Electron deficient – form with fewer than 8
electrons around atom
• Be, B
• Rare
Tend to form coordinate covalent bonds –
both electrons in shared pair donated by
single atom
+

Exceptions to Octet Rule
BeH2 – 4 electrons
BF3 – 6
electrons
Exceptions to Octet Rule
More than 8 valence electrons =
expanded octet
PCl5 SF6
d orbitals involved
• Only can occur for period 3 and higher,
not periods 1 or 2
Practice (Octet Exceptions)
Problems 47 - 49 page 260
Problems 102 (a-d), 104(a-d) page 273
Problem 7, page 980
Chapter 8 – Covalent Bonding
8.1
8.2
8.3
8.4
8.5
The Covalent Bond
Naming Molecules
Molecular Structures
Molecular Shape
Electronegativity and Polarity
Section 8.2 Naming Molecules
Specific rules are used when naming
binary molecular compounds, binary
acids, and oxyacids.
• Translate molecular formulas into binary molecular
compound names and also the reverse process.
• Name acidic solutions
Naming Binary Covalent Compounds
First element named first, using entire
element name
Second element named using same
procedure as for ionic compounds –
root of element name + ide ending
Use prefixes except if first element = 1
• Drop final letter in prefix if precedes
vowel
• Carbon monoxide , not monooxide
Prefixes in Covalent Compounds
Table 9-1, page 248
# Atoms
Prefix
# Atoms
Prefix
1
mono-
6
hexa-
2
di-
7
hepta-
3
tri-
8
octa-
4
tetra-
9
nona-
5
penta-
10
deca-
Naming Binary Covalent Compounds
Name of AlCl3 ?
Aluminum chloride
Name of PCl3 ?
Phosphorus trichloride
Name of Al2O3?
Aluminum oxide
Name of P2O5 ?
Diphosphorus pentoxide
The naming systems for ionic and covalent
compounds are different!!!
Common Names
Table 9-2, page 249
Form
ula
Common
Name
Molecular Compound
Name
H2O
Water
Dihydrogen monoxide
NH3
Ammonia
Nitrogen trihydride
N2H4
Hydrazine
Dinitrogen tetrahydride
N2O
Nitrous oxide Dinitrogen monoxide
(laughing gas)
NO
Nitric oxide
Nitrogen monoxide
Naming Acids
For our purposes, acids are what result
when molecules dissolved in water
produce H+ (hydrogen ions)
• HCl(g) in water  H+(aq) + Cl-(aq)
• Product is hydrochloric acid
Two common types
• Binary – H and one other element
• Oxyacid – H and an oxyanion
Naming True Binary Acids
Use prefix hydro- to name hydrogen
part of compound
For remainder, use a “form of the root”
of 2d element plus suffix –ic followed by
word acid
HCl – hydrochloric acid
H2S – hydrosulfuric acid
• Root of S for acid name not “sulf” as in
Na2S (sodium sulfide)
Naming Acids Similar
to Binary Acids (Rare)
If second part of compound is a
polyatomic anion that does not contain
oxygen (rare), use same system as for
a true binary acid employing the root
name for the anion
CN- – cyanide anion
HCN – hydrocyanic acid
Naming Oxyacids
Name is based solely on the anion
“A form of the root name of the anion” +
suffix + acid
Anion suffix
Acid Suffix
-ate
-ic
-ite
-ous
HNO3 Nitric acid
NO3- = nitrate
HNO2 Nitrous acid NO2- = nitrite
Naming Molecular Compounds
Flow Chart, Fig 9-9, page 251
Naming Molecular Compounds
Flow Chart, Fig 9-9, page 251
Acidic
Not Acidic
Practice
Problems 13-17 page 249 (binary
covalent)
Problems 18-22 page 250 (acids)
Problems 27-29 page 251 (mixed)
Problems 94-96(all a-d) page 273
Problems 97-98(all a-d) page 273
Problems 2 (a-f) page 874 (binary cov)
Problem 3 page 875 (acids)
Chapter 8 – Covalent Bonding
8.1
8.2
8.3
8.4
8.5
The Covalent Bond – Bond Strength
Naming Molecules
Molecular Structures
Molecular Shape
Electronegativity and Polarity
Section 8.1 The Covalent Bond
• Relate the strength of a covalent bond to its bond
length, bond order, and bond dissociation energy.
• Describe how the overall energy of a reaction (i.e.,
whether it is an endo- or exothermic reaction) is
related to the bond energies of the reactant and
product molecules.
Potential Energy (kJ/mol)
Covalent Bonding – Energy for H2
-432 kJ/mol
100
200
Internuclear Distance (pm)
Bond Strength & Bond Length/Order
Distance between bonding nuclei at
position of max attraction = bond length
Scale of bond length: ~10-10 m =100 pm
Bond order: Single 1 Double 2 Triple 3
Bond Strength & Bond Length/Order
Strength of bond related to bond length
Bond dissociation energy = energy needed
to break bond
Triple bond > double bond > single bond
Molecule Bond Length Dissoc. Energy
(pm)
kJ/mol
F2
143
159
O2
121
498
N2
110
945
Bond Strength & Bond Length/Order
Reaction Energies & Bond Energies
Chemical reaction
Bonds in reactant molecules broken
New bonds formed in product
molecules
CH4 + 2O2  2H2O + CO2
Breaking C-H bonds and O=O bonds
Making O-H bonds and C=O bonds
Reaction Energies & Bond Energies
CH4(g) + 2O2(g)  2H2O(g) + CO2(g)
Total energy change determined by
difference of energy of bonds broken
(reactant side) and formed (product
side)
• Endothermic – need more energy to break
than get back in formation
• Exothermic – bond formation energy
larger than energy needed to break bonds
Reaction Energies & Bond Energies
Enthalpy
-SBE
(products)
SBE
(reactants)
SBE
-SBE
(products)
(reactants)
BE = Bond energy
Chapter 8 – Covalent Bonding
8.1
8.2
8.3
8.4
8.5
The Covalent Bond
Naming Molecules
Molecular Structures
Molecular Shape
Electronegativity and Polarity
Section 8.4 Molecular Shapes
The VSEPR model is used to determine
molecular shape.
• Summarize the VSEPR bonding theory, including the
role of bonding and nonbonding pairs of electrons.
• Predict the shape of, and the bond angles in, a
molecule using VSEPR theory.
2 Simple Theories Related to Covalent
Bonding
Valence Shell Electron Pair Repulsion
Theory (VSEPR)
• Use Lewis structures to predict shape
Valence Bond Theory
• Extends Lewis bonding model to focus
on orbitals, particularly hybridized orbitals
VSEPR
Valence Shell Electron Pair Repulsion
Theory - allows us to predict geometry
Lewis structures tell us how the atoms
are connected to each other
Lewis structures don’t tell us anything
about shape
Shape of a molecule can greatly affect
its properties
Molecular Shape & Biological Sensors
For some biological systems, a
response is generated or a chemical
change is initiated when a molecular
key fits into correspondingly shaped
molecular lock
• Key is typically small molecule
• Lock is typically large molecule with a
shaped receptor site
 Only interacts with key of a specific
shape
Lewis Structure (a) & Tetrahedral
Geometry (b) for Methane (CH4)
VSEPR
Molecules take a shape that puts electron
pairs as far away from each other as
possible (electron pair repulsion)
Have to draw the Lewis structure to
determine categories of electron pairs
• bonding
• nonbonding lone pair
Lone pair take more space
Multiple bonds count as one pair
Balloon Analogy for the Mutual
Repulsion of Electron Groups
Two
Three
Four
Five
Six
Number of Electron Groups
VSEPR
The number of pairs determines
• bond angles
• underlying structure
The number and position of atoms
determines
• actual molecular shape
VSEPR – Underlying Shapes
# Elec. pairs Bond Angles Shape
2
180°
Linear
120° Trigonal Planar
3
4
5
6
109.5°
90° &
120°
90°
Tetrahedral
Trigonal
Bipyramidal
Octahe
dral
Actual Molecular Shapes
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
2
3
3
4
4
4
2
3
2
4
3
2
0
0
1
0
1
2
linear
trigonal planar
bent
tetrahedral
trigonal pyramidal
bent
Actual Molecular Shapes
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
5
5
5
5
5
4
3
2
0
1
2
3
trigonal bipyrimidal
See-saw
T-shaped
linear
Actual Molecular Shapes
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
6
6
6
6
6
6
5
4
3
2
0
1
2
3
4
Octahedral
Square Pyramidal
Square Planar
T-shaped
linear
Relative Sizes: Bonding Pairs vs Lone Pairs
CH4
NH3
H2O
Molecular Geometry
Can predict geometry around each
atom center and build overall molecular
geometry piece by piece
N
O1
Glycine
C2
C1
O2
Chapter 8 – Covalent Bonding
8.1
8.2
8.3
8.4
The Covalent Bond
Naming Molecules
Molecular Structures
Molecular Shape (extension of book)
Valence Bond Theory - Orbital Overlap
Hybrid Orbitals
Quantum mechanical calculations
8.5 Electronegativity and Polarity
Section 8.4 Molecular Shapes
The Valence Bond model is used to
determine molecular shape via the
concept of overlap of orbitals,
particularly hybrid orbitals.
• Describe the valence bond model of bonding
• Explain the similarities and differences between the
Lewis and valence bond models of chemical bonds.
• Describe sigma and pi bonds and identify these
bonds within molecules.
• Define hybridization.
Section 8.4 Molecular Shapes
• Relate the type of hybridization (sp3, sp2, etc.) to the
VSEPR geometry of a molecule
• Identify the specific type of hybridization that occurs
within a given molecule and identify the specific
orbitals (hybrid or non-hybrid) that are involved in
each sigma and pi bond.
• Explain how quantum mechanics and the wave
function concept can be applied to a molecule.
Valence Bond Theory
Lewis structures indicate status of electrons
• Shared in bond
• Lone pair
No information about orbitals involved
Valence bond theory
• Bonds are formed by overlap of half-filled
atomic orbitals
• Orbital geometry can give direct information
about molecular shape
Sigma Bonds
Single covalent bonds = sigma bond
• Symbol Greek letter 
Occurs when electron pair shared in
area centered between two atoms
Atomic orbitals overlap end to end,
forming a bonding orbital
• Localized region where bonding
electrons will most likely be found
Sigma Bond Formation by
Orbital Overlap
Two s orbitals
overlap
Sigma Bond Formation
H2
HF
F2
Two s
orbitals
overlap
Two p
orbitals
overlap
Sigma Bonding – F2


px1py2pz2
 
—F
F
 


Area of overlap
for atomic p 1p 2p 2
x y z
orbitals
Pi Bond ()
Formed when parallel orbitals
overlap to share electrons
Shared pair occupies space
above and below a line
connecting atoms
Multiple bonds always have
one sigma and at least one pi
bond
• Double: 1 , 1  bond
• Triple: 1 , 2  bonds
Sigma & Pi Bonding
Sigma () and Pi () Bonds
Hybrid Orbitals
For correct geometry of polyatomic
molecules using the valence bond
model, have to use concept of hybrid
orbitals
• CH4 has 109.5 angles, but atomic p
orbitals are at right angles to each other
Hybrid Orbitals
Hybrid orbitals – orbitals obtained when 2 or
more nonequivalent orbitals combine to
form an equal number of identical,
degenerate orbitals
Hybridization – mixing of atomic orbitals in
an atom (usually a central atom) to
generate a set of hybrid orbitals
Use VSEPR logic to determine geometry of
hybrid orbitals formed
Valence
Orbitals on a
Free Carbon
Atom:
2s, 2px, 2py,
and 2pz
s
py
px
pz
Formation of sp3 Hybrid Orbitals From
Original Valence Orbitals
Hybrid
ization
Cross
Section
of sp3
Orbital
Energy-Level Diagram Showing
Formation of Four sp3 Orbitals
Hybridization
Orbitals in
free C atom
C Orbitals in
CH4 molecule
Valence Bond Theory Treatment of CH4
2p
2s
C
C*
sp3
C* (sp3)
H
4H
1s
1s
1s
1s
C H
H
H
Overlap of sp3 hybrid orbitals on C with 1s
orbitals on H atoms gives 4 C-H (sp3)-1s  bonds
oriented 109.47° from each other
Has tetrahedral geometry predicted by VSEPR
Tetrahedral
Set of Four
sp3 Orbitals
Forming
Sigma
Bonds with
s Orbitals of
Four
Hydrogen
Atoms
(CH4)
Formation of sp2 Hybrid Orbitals from s,
px, and py Atomic Orbitals
Hybrid
ization
Energy-Level Diagram Showing
Formation of Three sp2 Orbitals
Hybridization
Orbitals in
free C atom
Orbitals in sp2
hybridized C
Note: Inconsistent with actual bonding – 4
valence electrons populate only sp2 orbitals
(Aufbau) leaving only 1 unpaired electron in sp2
An sp2 Hybridized C Atom
Formation of sp Hybrid Orbitals
from s and px, Atomic Orbitals
Hybrid
ization
Energy-Level Diagram Showing
Formation of Two sp Hybrid Orbitals
Hybridization
Orbitals in
free C atom
Orbitals in sp
hybridized C
Note: Inconsistent with actual bonding – 4
valence electrons should populate only sp orbitals
(Aufbau) leaving no unpaired electrons
Orbitals of sp Hybridized Carbon Atom
sp3d (dsp3) Hybrid Orbitals
Can only occur for periods 3 & higher (need
d orbitals) – example shown is for P
Linked to geometry with 5 pairs (trigonal
bipyramid)
3s
3p
P
3d
P*
P* (sp3d)
3s
3pz
3py
3px
3d
3dz2
sp3dz2
Set of dsp3 Hybrid Orbitals on a
Phosphorus Atom
sp3d2 (d2sp3) Hybrid Orbitals
Can only occur for periods 3 & higher (need
d orbitals)
Linked to geometry with 6 pairs (octahedral)
Example on next slide for S
S - Octahedral Set of d2sp3 Orbitals
Relationship among the number of
effective pairs, geometry, and the
hybrid orbital set required to obtain this
geometry shown on the following two
slides
#
Geometry
Hybridization
2
Linear
sp
3
Trigonal sp2
planar
4
Tetra
hedral
sp3
#
Geometry
Hybridization
5
Trigonal sp3d
bipyramidal
6
Octa
sp3d2
hedral
Geometry & Hybridization - Steps
1. Draw Lewis structure
2. Determine # of effective electron pairs
(count double & triple bonds as one pair)
3. Determine basic geometry from number of
pairs (e.g., 5 pairs = trigonal bipyramid)
4. Determine hybridization type from number
of pairs (e.g., 5 pairs = sp3d)
5. Form single (sigma) bonds from hybrid
orbitals; lone pairs also go in hybrid orbitals
6. Form pi bonds using unhybridized orbitals
Geometry & Hybridization - Steps
Following slides give examples of using the
steps listed on previous slide for these
molecules:
1. Ammonia
2. Ethylene
3. Diatomic nitrogen
4. Acetylene
5. Carbon dioxide
6. Phosphorus pentachloride
N in Ammonia
sp3 Hybridized (4 pairs)
N in Ammonia
Trigonal
pyramidal
molecule with
lone pair
occupying
hybrid orbital
Sigma & Pi Bonds Using Hybrid
Orbitals - Ethylene
Three electron pairs for C  sp2
hybridization & trigonal planar geometry
C 1s22s22p2  1s22(sp2)32p1
Hydrogens have 1s1 orbitals (spherical)
 Bonds in Ethylene – Top View
Sigma () bonds
Sigma and Pi Bonds in Ethylene
Pi () bond
Sigma Bonds in Ethylene
Because each C has trigonal planar
geometry, entire molecule is planar
N2 Bonding
lone pair
sigma
lone pair
sp hybridized
(2 pairs)
py
sp
One sigma bond
sp
pz
Two sp hybrid
orbitals and two
normal p orbitals
Two pi bonds
Sigma and Pi Bonds in Acetylene
Two electron pairs for C  sp hybridization,
linear geometry (triple bond = single pair)
C 1s22s22p2  1s22(sp)22p2
Hydrogens have 1s1 orbitals (spherical)
Orbitals of sp Hybridized Carbon Atom
Sigma and Pi Bonds in Acetylene
sp hybrid orbitals on C form single
(sigma) bond with H and other C
Remaining two unhybridized p orbitals
overlap to form two pi bonds
Sigma and Pi Bonds in Acetylene
Sigma & Pi Bonds Using Hybrid
Orbitals in CO2
Two electron pairs for C  sp hybridization,
linear geometry (double bond = single pair)
C 1s22s22p2  1s22(sp)22p2
Three electron pairs for O  sp2
hybridization & trigonal planar geometry
O 1s22s22p4  1s22(sp2)5p1
Orbitals of sp Hybridized Carbon Atom
Orbital Arrangement for an sp2
Hybridized Oxygen Atom
Sigma Bonds using Hybrid Orbitals in
CO2 Molecule
Sigma () bonds
Sigma and Pi Bonds Using Hybrid
Orbitals in CO2
Sigma Bonds Using Hybrid Orbitals in
PCl5
Five electron pairs for P  sp3d hybridization
& trigonal bipyramidal geometry
P [Ne]3s23p3  [Ne](sp3d)5
Four electron pairs for Cl  sp3 hybridization
& tetrahedral geometry
Cl [Ne]3s23p5  [Ne]3(sp3)7
Set of dsp3 Hybrid Orbitals on a
Phosphorus Atom
Structure of PCI5 and Orbitals Used to
Form Sigma Bonds
Sigma () bond
Lone
pairs on
Cl in sp3
orbitals
Geometry & Hybridization
Supply for each indicated atom in structure
# of sigma & pi bonds in molecule?
H
O
H
C O C
C C
H
H
C
N
12 , 4 
33pairs
pairs
24 pairs
4
pairs
H Trigonal
(2
lone)
Trigonal
Linear
Tetrahedral
planar
planar
Bent
3
sp
sp
2
sp
2
3
sp
sp
Practice
(Shape, Angles, Hybridization)
Problems 56 – 60, 65 - 67 page 264
Problems 108,110 - 112 page 275
Problem 8 page 980
Quantum Mechanics & Molecules
Quantum Mechanics & Molecules
Y (wave function) exists for entire
molecule and can be obtained from
solution to Schrodinger wave equation
written for the molecule
Y2 - Square of Y gives probability of
finding electron at particular position
around molecule – defines what is
called a molecular orbital (MO)
Quantum Mechanics & Molecules
Using certain types of approximations and
today’s computers, wave functions for
molecules (not individual atoms) can be
obtained and molecular properties
calculated from this information
Energy, absorption spectrum, dipole
moment, etc
Molecular orbital theory is most advanced
way of describing covalent bonding
Chapter 8 – Covalent Bonding
8.1
8.2
8.3
8.4
8.5
The Covalent Bond
Naming Molecules
Molecular Structures
Molecular Shape
Electronegativity and Polarity
Section 8.5 Electronegativity and Polarity
A chemical bond’s character is related
to each atom’s attraction for the
electrons in the bond.
• Describe how electronegativity is used to determine
bond type and characterize bonds between given
pairs of atoms as being polar or nonpolar.
• Compare and contrast polar and nonpolar covalent
bonds and polar and nonpolar molecules.
• Describe the term “dipole moment” and relate it to the
terms polar and nonpolar.
Section 8.5 Electronegativity and Polarity
• Identify molecules as being polar or nonpolar.
• Describe how polarity affects the solubility of one
substance in another substance.
• Describe how polarity can give rise to intermolecular
forces.
Polar Covalent Bonds
Polarity of bond determined by
electronegativity difference
Difference = 0 Nonpolar
Difference > 0 Polar
Very large differences
• No longer covalent compound
75 100
Ionic Bonds
50
25
% Ionic Character
EN Difference & Bond Character
0
Covalent
Bonds
1.0
2.0
3.0
Electronegativity Difference
Relationship Between EN
Difference and Bond Type
Relationship Between EN
Difference and Bond Type
EN=0
Nonpolar
Covalent
EN=medium EN=large
Polar
Ionic
Covalent
Scalars & Vectors
Scalar
• Completely specified by magnitude
and units
Vector
• Has magnitude, direction, and units
v = 3.5 m/s
(scalar)
v = 3.5 m/s to northeast (vector)
Trigonometric Functions
Pythagorean Theorem
Dipole Moment
Two equal and opposite charges +Q and -Q
separated by a distance l have a dipole
r
moment p:
p =  p = Q  l
(vector points from –Q to +Q)
+Q
r
p
l
-Q
Polarity and Dipole Moment
+
Dipole
_
Dipole moment is a vector pointing from
center of - charge to center of + charge
Magnitude proportional to size of charges
and to separation distance
All polar covalent bonds have a dipole
moment
Polarity and Dipole Moment
p =  p = Q  l +Q
r
-Q
p
l
Units of p are Debye units (D)
% ionic character of bond determined by
size of measured dipole moment relative to
value calculated from using full (ionic)
charges as Q
Dipole Moments of Gas Phase
Molecules
Dipole Moment
Dipole
moments from
bonds add as
vectors to give
dipole moment
of molecule
r
r
r
p  p1 + p 2
(net dipole moment)
Molecular Polarity – Linear Molecule
O=C bond polar; bonding electrons pulled
equally toward both O ends of molecule
Net result is nonpolar molecule (dipole
moments of bonds cancel each other)
[note: red arrows are opposite dipole direction]
Molecular Polarity – Bent Molecule
H-O bond polar
Both sets of bonding electrons pulled toward O
end; net result is polar molecule (y components
of bond dipole moment add, x components
cancel)
[note: red arrows are opposite dipole direction]
Polar Molecules
Molecule can have polar bonds but be
a nonpolar molecule
-
-
+
-
+
Polar
+
-
-
Nonpolar
Polar Bonds in Nonpolar Molecules
In symmetric molecules, vector addition
of bond dipole moments results in zero
dipole moment for the molecule
All molecules having basic VSEPR
shapes & equal bonds are nonpolar
-
+
Linear
Trigonal
- Planar +
Practice
(Polar Bonds & Polar Molecules)
Problems 74 – 77 page 270
Problems 117 – 123 page 275
Problem 9 page 980
Polarity Effects
Polarity of molecule determines
solubility characteristics – “like
dissolves like”
Oil (nonpolar)
and water
(polar) don’t
mix
Dipole in an Electric Field
The + and – charges in an electric
dipole are pulled in opposite directions
in an electric field, producing a net
torque on the dipole, and orienting it.
Dipole in Electric Field – HF Molecule
+δ
-δ
Field
Off
F
H
+
Field
On
-
Polar Molecule & Electric Field
Polar molecules affected by electric
field in an EM wave
Oscillating field twists water molecule
and energy transferred (heats up)
Basis for microwave oven operation
Properties of Covalent Compounds
Bonding types affect properties
Many properties controlled by
intermolecular forces
• Forces between molecules
• Also known as van der Waals forces
Intermolecular forces are weaker than
chemical bonds
[Note: intermolecular forces treated in more
depth in section 12.2 – Forces of Attraction]
Intermolecular Forces
Forces between nonpolar molecules
relatively weak
• Tend to be gases or volatile liquids
• O2, N2, small hydrocarbons
Forces between polar molecules are
stronger due to dipole-dipole forces
• Hydrogen bonding a particular strong
version - H and F, O, or N
Hydrogen Bonding – Water Pentamer
Hydrogen
Bonds
Hydrogen Bonding in Nylon
Hydrogen bonding helps make nylon strong
End of Chapter 8