Chapter 8: Covalent Compounds
Polarity
A molecule, such as HF, that has a
center of positive charge and a
center of negative charge is said to
be polar, or to have a dipole
moment.
Polarity of Molecules
Dipole: def: two charges, equal in magnitude and
opposite in sign, are separated by a distance
Polarity of Polyatomic Molecules
•Each bond can be polar.
•The orientation of these polar bonds determines
whether the molecule is polar overall.
•It is possible for a molecule with polar bonds to be
either polar or non-polar.
Polarity of Molecules
Dipole Moments of Polyatomic Molecules
Example: in CO2, each C-O dipole is canceled because
the molecule is linear. In H2O, the H-O dipoles do not
cancel because the molecule is bent.
Polarity of Molecules
Dipole Moments of Polyatomic Molecules
Models

Models are attempts to explain how nature operates
on the microscopic level based on experiences in the
macroscopic world.

Properties of Models:

A model does not equal reality.

Models are oversimplifications, and are therefore
often wrong.

Models become more complicated as they age.

We must understand the underlying assumptions
in a model so that we don’t misuse it.
Lewis Structure



Shows how valence electrons are
arranged among atoms in a
molecule.
Valence electrons are found in the
outermost energy level of an
atom. They are involved in
bonding.
Reflects the central idea that the
stability of a compound relates to
noble gas electron configuration.
Lewis Dot Structure Rules:







Treat ions separately (e.g. NH4Cl)
Count only valence electrons
Assemble bonding framework
Fill up non-bonding electrons on outer
atoms
Fill up non-bonding electrons on inner
atoms
Calculate Formal Charge
Minimize Formal Charge
To Complete a Lewis Structure:

Must be able to recognize
polyatomic ions


Must be able to identify
valence electrons
Must be able to construct
Bond framework
Hints on Lewis Dot Structures
1.
2.
3.
4.
5.
6.
Octet “rule” is the most useful guideline.
Carbon always forms 4 bonds.
Other 2nd row elements (N, O, F) observe
the octet rule.
Hydrogen forms one bond with other
atoms to complete its octet with 2
electrons.
2nd row elements B and Be often have
fewer than 8 electrons around themselves
- they are very reactive.
3rd row and heavier elements CAN exceed
the octet rule using empty valence d
orbitals.
Notes About the Octet Rule
7.
8.
9.
10.
11.
When multiple bonds are forming, they are
usually between C, N, O or S.
Nonmetals can form single, double, and
triple bonds, but not quadruple bonds.
Always account for single bonds and lone
pairs before forming multiple bonds.
Look for resonance structures.
When writing Lewis structures, satisfy
octets first, then place electrons around
central atoms having available d orbitals.
5+(3*7)=26 e-
PCl3
Bonding Pairs
Lone Pairs
(a.k.a. nonbonding electrons)
Try Some Examples:
C2H5OH
 CH3CH2NH2


Cl2CO

Ozone (O3)

NO2
Formal Charge
Difference between the # of valence electrons in the free atom
and the # of electrons assigned to that atom in the Lewis
structure.
1

FC  G.N. -  # BE  # LPE 
2

FC = formal charge; G.N. = Group Number
#BE = bonding electrons; #LPE = lone pair electrons
If Step 4 leads to a positive formal charge on an inner atom beyond
the second row, shift electrons to make double or triple bonds to
minimize formal charge, even if this gives an inner atom with
more than an octet of electrons.
Formal Charge
O C O
(-1)
(0)
(+1)
Not as good
O C O
(0)
(0)
(0)
Better
Covalent Bonding
Multiple Bonds
• It is possible for more than one pair of electrons to
be shared between two atoms (multiple bonds):
• One shared pair of electrons = single bond (e.g. H2);
• Two shared pairs of electrons = double bond (e.g. O2);
• Three shared pairs of electrons = triple bond (e.g. N2).
H H
O O
N N
Octet in each case
• Generally, bond distances shorten with multiple
bonding.
Resonance

Occurs when more than one valid
Lewis structure can be written for a
particular molecule.
Odd Number of Electrons…
NO
Number of valence electrons = 11
N O
N O
Resonance occurs when more than one valid Lewis structure can be
written for a particular molecule (i.e. rearrange electrons)
NO2
Number of valence electrons = 17
O N O
O N O
O N O
Molecules and atoms which are neutral (contain no formal charge) and with an
unpaired electron are called Radicals
Beyond the Octet


Elements in the 3rd period or higher
can have more than an octet if
needed.
Atoms of these elements have
valence d orbitals, which allow them
to accommodate more than eight
electrons.
More than an Octet…
Elements from the 3rd period and beyond, have ns, np and
unfilled nd orbitals which can be used in bonding
Cl
PCl5
P : (Ne) 3s2 3p3 3d0
Number of valence electrons = 5 + (5 x 7) = 40
Cl
P
Cl
SF4
Cl
Cl
S : (Ne) 3s2 3p4 3d0
Number of valence electrons = 6 + (4 x 7) = 34
F
The Larger the central atom, the more atoms
you can bond to it – usually small atoms such
as F, Cl and O allow central atoms such as P
and S to expand their valency.
F
S
F
F
Less than an Octet…
BCl3
Cl
B
Cl
Cl
Group 13 atom only has six electrons
around it
Molecular Shapes
• Lewis structures give atomic connectivity:
they tell us which atoms are physically
connected together. They do not tell us the
shape.
• The shape of a molecule is determined by its
bond angles.
• Consider CCl4: experimentally we find all ClC-Cl bond angles are 109.5.
Therefore, the molecule cannot be planar.
All Cl atoms are located at the vertices of a
tetrahedron with the C at its center.
Molecular Shape of CCl4
VSEPR Theory
 In order to predict molecular shape, we
assume the valence electrons repel each other.
Therefore, the molecule adopts whichever 3D
geometry minimized this repulsion.
 We call this process Valence Shell Electron
Pair Repulsion (VSEPR) theory.
Why is VSEPR Theory Important?
 Gives
a specific shape due to the
number of bonded and nonbonded electron pairs in a
molecule
 Tells us the actual 3-D structure
of a molecule
 In bonding, electron pairs want to
be as far away from each other as
possible.
VSEPR and Resulting Geometries
How does VSEPR THEORY work?
We can use VSEPR theory using 4
steps.
1. Draw the Lewis Structure for the
molecule.
Example: SiF4 F
F-Si-F
F
How does VSEPR THEORY work?
We can use VSEPR theory using 4
steps
1. Draw the Lewis Structure for the
molecule.
2. Tally the number of bonding
pairs and lone (non-bonding)
pairs on the center atom.
F
F-Si-F
F
Bonding pairs: 4
Lone pairs on central atom: 0
How does VSEPR THEORY work?
We can use VSEPR theory
using 4 steps
1. Draw the Lewis Structure
for the molecule
2. Tally the number of
bonding pairs and lone
pairs on the center atom.
3. Arrange the rest of the
F
atoms so that they are as
Si
far away from each other F
F
as possible.
F
How does VSEPR THEORY work?
We can use VSEPR theory using 4 steps
1. Draw the Lewis Structure for the
molecule
2. Tally the number of bonding pairs
and lone pairs on the center atom.
3. Arrange the rest of the atoms so that
they are as far away from each other
as possible
4. Give the type of geometry the
molecule has:
Tetrahedral
Another Example:
To determine the electron pair geometry:
1) draw the Lewis structure;
2) count the total number of electron pairs around the
central atom.
3) arrange the electron pairs in one of the geometries
to minimize e--e- repulsion.
4) multiple bonds count as one bonding pair for VSEPR
The VSEPR Model
Predicting Molecular Geometries
The VSEPR Model
Predicting Molecular Geometries
The VSEPR Model
Difference between geometry and shape
Geometry:
We determine the geometry only looking at electrons.
All the atoms that obey the octet rule have the same
tetrahedral-like geometry.
Shape:
We name the shape by the positions of atoms.
We ignore lone pairs in the shape.
The VSEPR Model
Predicting Shape
Shape
The VSEPR Model
Predicting Shape
Shape
The VSEPR Model
The Effect of Nonbonding Electrons and
Multiple Bonds on Bond Angles
By experiment, the H-X-H bond angle decreases on
moving from C to N to O:
H
H C H
H
109.5O
H N H
H
107O
O
H
H
104.5O
Since electrons in a bond are attracted by two nuclei, they do
not repel as much as lone pairs.
Therefore, the bond angle decreases as the number of lone
pairs increase.
The VSEPR Model
The Effect of Nonbonding Electrons and
Multiple Bonds on Bond Angles
Similarly, electrons in multiple bonds repel more than
electrons in single bonds.
Cl
111.4o
Cl
C O
124.3o
The VSEPR Model
Molecules with Expanded Valence Shells
Atoms that have expanded octets have AB5 (trigonal
bipyramidal) or AB6 (octahedral) electron pair
geometries.
Examples:
PF5 trigonal bipyramidal
SCl6 octahedral
The VSEPR Model
Molecules with Expanded Valence Shells
The VSEPR Model
Molecules with Expanded Valence Shells
The VSEPR Model
Molecules with More than One Central Atom
In acetic acid, CH3COOH, there are three central
atoms.
We assign the geometry about each central atom
separately.
Hybrid Orbitals
 In
bonding, s and p orbitals are
used in bonding. It is easy to tell
which ones are used by looking at
our molecule.
 For example, CH4. Looking again
at the Lewis structure, we see
that there are 4 bonds. We call
this sp3 hybridized.
Hybrid Orbitals
 Regions
of electron density-EACH
BOND AND LONE PAIR OF
ELECTRONS ON THE CENTRAL
ATOM IS KNOWN AS A REGION OF
ELECTRON DENSITY.
 2 regions of electron density-sp
hybridized
 3 regions of electron density-sp2
hybridized
 4 regions of electron density-sp3
hybridized
Hybridization
sp Hybrid Orbitals
The two lobes of an sp hybrid orbital are 180 apart.
Hybrid Orbitals
sp2 Hybrid Orbitals
Important: when we mix n atomic orbitals we must get
n hybrid orbitals.
sp2 hybrid orbitals are formed with one s and two p
orbitals. (Therefore, there is one unhybridized p orbital
remaining.)
The large lobes of sp2 hybrids lie in a trigonal plane.
All molecules with trigonal planar electron pair
geometries have sp2 orbitals on the central atom.
Hybridization
Hybridization
sp3 Hybrid Orbitals
sp3 Hybrid orbitals are formed from one s and three p
orbitals. Therefore, there are four large lobes.
Each lobe points towards the vertex of a tetrahedron.
The angle between the large lobes is 109.5
All molecules with tetrahedral electron pair geometries
are sp3 hybridized.
Hybridization
Hybrid Orbitals
Hybrid Orbitals
Hybrid Orbitals
Summary
To assign hybridization:
1.
2.
3.
4.
Draw a Lewis structure.
Assign the geometry using VSEPR theory.
Use the geometry to determine the hybridization.
Name the shape by the positions of the atoms.
Hybridization and Multiple Bonds
•Multiple bonds overlap differently and are
called s-bonds and p-bonds
•All single bonds are s
•Double bonds contain
1 s and 1 p bond
•Triple bonds contain
1 s and 2 p bonds
Bond Energy
Covalent Bonding & Orbital Overlap
•As two nuclei approach each other their
atomic orbitals overlap.
•As the amount of overlap increases, the
energy of the interaction decreases.
•At some distance the minimum energy is
reached.
•The minimum energy corresponds to the
bonding distance (or bond length).
Covalent Bonding & Orbital Overlap
•As the two atoms get closer, their nuclei
begin to repel and the energy increases.
•At the bonding distance, the attractive
forces between nuclei and electrons just
balance the repulsive forces (nucleus-nucleus,
electron-electron).
Bond Energies
Bond breaking requires energy
(endothermic).
 Bond formation releases energy
(exothermic).
H = D(bonds broken) - D(bonds formed)

energy required
energy released
Bond Energies