Covalent Bonding
Lesson 1:Covalent Bonding
 Covalent bonds: atoms held together by sharing




electrons.
Molecules: neutral group of atoms joined together by
covalent bonds.
Diatomic molecule: molecule consisting of 2 atoms.
Remember them: F2, Cl2, I2, Br2, H2, N2, O2
Molecules tend to have lower melting and boiling
points than ionic compounds.
YouTube - Making Molecules with Atoms
Molecular Formula
 Shows how many atoms of each element a molecule
contains.
 Naming binary molecular compounds
 Composed of two nonmetals; often combine in more
than one way. Ex. CO and CO2
 Prefixes are used to name binary molecular compounds.
Prefix
Mono-
Di-
Tri-
Tetra-
Penta-
Hexa-
Hepta-
Octa-
Nona-
Deca-
Number
1
2
3
4
5
6
7
8
9
10
Binary Compounds
Containing Two Nonmetals
To name these compounds:
1) give the name of the less electronegative element first
with the Greek prefix indicating the number of atoms
of that element present
2) After give the name of the more electronegative nonmetal with the Greek prefix indicating the number of
atoms of that element present and with its ending
replaced by the suffix –ide.
3) Do not use the prefix mono- if required for the first
element.
Binary Molecular Compounds
N2O
N2O3
N2O5
dinitrogen monoxide
dinitrogen trioxide
dinitrogen pentoxide
ICl
ICl3
iodine monochloride
iodine trichloride
SO2
sulfur dioxide
SO3
sulfur trioxide
YouTube - Naming molecular compounds
Binary Molecular Compounds
Containing Two Nonmetals
As2S3
1. ________________
SO2
2. ________________
diarsenic trisulfide
sulfur dioxide
P2O5
diphosphorus pentoxide
____________________
CO2
4. ________________
carbon dioxide
3.
5.
N2O5
dinitrogen pentoxide
____________________
6.
H2O
dihydrogen monoxide
____________________
Naming Binary Compounds
Binary Compound?
Yes
Metal Present?
No
Molecule
Use Greek
Prefixes
Yes
Does the metal form
more than one cation?
No
Ionic compound (cation has
one charge only)
Use the element
name for the cation.
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 98
Yes
Ionic compound (cation has more
than one charge) Determine the
Charge of the cation; use a Roman
numeral after the cation name.
Classwork #1:
Do handout “Naming
Molecules”
Lesson 2: The Nature of Covalent
Bonding
 Introduction with balloon activity
 octet rule: electron sharing occurs usually so that
atoms attain the electron configurations of noble
gases.
 Single covalent bond: two atoms held together by
sharing a pair of electrons. Shown as two dots or as a
long dash.
 A pair of valence electrons that is not shared between
atoms is called an unshared pair.
O
H
H
H
O H
H
or
O
H
 Double bonds: covalent bond formed by sharing two
pairs of electrons
 Triple bonds: covalent bond formed by sharing three
pairs of electrons.
Covalent bonding with equal sharing
of electrons occurs in diatomic
molecules formed from one element.
hydrogen
chlorine
iodine
nitrogen
A dash may replace a pair of dots.
Classwork: introduction to lewis
structures.
Lesson 3:Molecular Structure
 Structural formula: uses symbols and bonds to show
relative position of atoms.
 Steps to determine Lewis structures for molecules
1. Predict the location of certain atoms.
 Hydrogen is always an end atom
 The least electronegative atom is the central atom
(usually the one closer to the left on periodic table)
2. Find the total number of electrons available for bonding.
(# of valence electrons of atoms in molecule)
3. Determine the number of bonding pairs by dividing the
total number of electron by 2
4. Place one bonding pair (single bond) between central
atom and terminal atoms.
5. Subtract pairs used in step 4 from bonding pairs in step
3. Place lone pairs around each terminal atom bonded to
the central atom to satisfy the octet rule. Any remaining
pairs are assigned to the central atom.
6. If the central atom does not have an octet, convert one or
two of the lone pairs on the terminal atoms to a double or
a triple bond between central and terminal atom. Some
elements like Be, B, Al do not form a complete octet, S and
P can have more than 8 valence electrons.
Ex. 1 Draw the lewis structure for ammonia, NH3
Hydrogen is an end atom and nitrogen is the central
atom.
2. Total number of valence electrons:
(1 nitrogen x 5 valence electrons)+ (3 hydrogens x 1 valence
electron)= 8 valence electrons.
3. Total number of bonding pairs= 8/2 = 4
4. Draw single bond from each H to N
1.
H
N
H
H
Ex. 1 Draw the lewis structure for ammonia, NH3
5. Subtract the number of pairs of electrons used from the
total pairs of electrons: 4-3 =1 pair available
One lone pair remains, hydrogen can have only one bond,
assign the lone pair to the central atom, N.
H
N
H
H
Ex. 2 Draw the lewis structure for carbon dioxide, CO2
Oxygen atoms are end atoms and carbon is the central
atom.
2. Total number of valence electrons:
(1 carbon x 4 valence electrons)+ (2 oxygen x 6 valence
electron)= 16 valence electrons.
3. Total number of bonding pairs= 16/2 = 8
4. Draw single bond from each C to O
1.
O
C
O
Ex. 1 Draw the lewis structure for carbon dioxide, CO2
5. Subtract the number of pairs of electrons used from the
total pairs of electrons: 8-2 =6 pair available
Add three pairs of electrons to each oxygen.
O
C
O
Ex. 1 Draw the lewis structure for carbon dioxide, CO2
6. No lone pairs remain for carbon. Carbon does not have
an octet, use a lone pair from each oxygen to form a
double bond with the carbon atom.
O
C
O
O
C
O
CW: lewis structures handout part 1
Lesson :4 Exception to octet rule
 Some molecules have an odd number of valence electrons
and cannot form an octet around each atom.
 Some molecules form with fewer than eight electrons
present around an atom. Ex. Boron
 Some compounds have central atoms with more than 8
electrons. This is called an expanded octet. Ex. S, Xe and P
Ex. 3 Draw the lewis structure for XeF4 (exception octet
rule)
F is an end atom and nitrogen is the central atom.
2. Total number of valence electrons:
(1 xenon x 8 valence electrons)+ (4 fluorines x 7 valence
electron)= 36 valence electrons.
3. Total number of bonding pairs= 36/2 = 18
4. Draw single bond from each F to Xe
1.
F
F
Xe F
F
Ex. 1 Draw the lewis structure for XeF4 (exception octet
rule)
5. Subtract the number of pairs of electrons used from the
total pairs of electrons: 18-4 =14 pairs available
14 lone pairs remain, place them around each fluorine so
that each fluorine has 8 valence electrons
F
F
Xe F
F
Ex. 1 Draw the lewis structure for XeF4 (exception octet
rule)
6. There are 2 pairs of electrons still available, place around
Xe which is capable of having more than 8 valence
electron.
F
F
Xe F
F
Molecular Shape
 VSEPR (Valence shell electron pair repulsion) Model
 The repulsion between electron pairs in a molecule
result in atoms existing at fixed angles from each other.
(Remember balloon activity)
 Shared electron pairs repel each other
 A greater repulsion occurs between unshared electron
pairs and shared electron pairs.
Bonding and Shape of Molecules:
Count number of bonds and unshared pairs of electrons
AROUND CENTRAL ATOM and then use table below to
determine shape of molecule.
Number
of Bonds
Number of
Unshared Pairs
0
3
0
4
0
3
1
2
2
Shape
Examples
-Be-
Linear
BeCl2
Trigonal planar
BF3
Tetrahedral
CH4, SiCl4
Pyramidal
NH3, PCl3
Bent
H2O, H2S, SCl2
B
C
:
2
Covalent
Structure
:
N
O:
Use table on last slide to determine
shape of molecule.
SO2
O
H
Shape: bent
2 bonds and 2
unshared pairs
H
Carbon tetrachloride
Cl
Cl C Cl
Cl
CCl4
Cl
C
Cl
109.5o
Cl
Cl
Shape: Tetrahedral
4 bonds and 0 unshared pairs.
Carbon tetrachloride – “carbon tet” had been used as dry cleaning solvent
because of its extreme non-polarity.
Classwork: do in your notebook
 Determine the shape for the following molecules (first
draw the lewis structure for the molecule and then use
the table on slide #7 to determine the shape taking in
consideration the number of bonds and unshared
pairs of electrons around the CENTRAL ATOM.)
1. BF3
2. OCl2
3. CF4
4. NH3
5. BeI2