Chapter 9
Covalent Bonding
9.1 The Covalent Bond
 Covalent Bond – A chemical bond that
results from the sharing of valence electrons.
– Between nonmetallic elements of similar
electronegativity.
 A group of atoms united by a covalent bond is
called a molecule.
 A substance made up of molecules is called a
molecular substance.
 Molecules may have as few as two atoms.
(CO), (O2)
Polar Covalent Bonds: Unevenly
matched, but willing to share.
 To describe the composition of a
molecular compound we use a
molecular formula.
 The molecular formula tells you how
many atoms are in a single molecule of
the compound.
– The molecular formula for oxygen is
O2, which tells you that this molecule
contains 2 atoms of oxygen.
 The molecular formula for sucrose
(table sugar) is C12H22O11. This
indicates that one sucrose molecule
contains 12 carbon atoms, 22 hydrogen
atoms and 11 oxygen atoms.
 The empirical formula shows the ratio
of atoms in the molecule. You get it by
dividing all subscripts by the smallest
subscript and rounding. C6H12O6
becomes CH2O.

Molecules are often shown using a
structural formula.
 A structural formula specifies which
atoms are bonded to each other in a
molecule.
– One type is the Lewis structures.
– The Lewis structure uses dots as
before and has a circle drawn
around the dots that represent
valence electrons that are joined in
a covalent bond.

The principle that describes covalent
bonding is the Octet rule.
– Atoms within a molecule will share
electrons so as to acquire a full set
(8) of valence electrons.

If each atom has 7 valence electrons,
than they will share one so that they
both will have 8. The sharing of that
electron is a Single covalent bond.
Single Covalent bonding
Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
 Both end with full orbitals

F
F
Covalent bonding
Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
 Both end with full orbitals

F F
8 Valence
electrons
Covalent bonding
Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
 Both end with full orbitals

8 Valence
electrons
F F




If more than two atoms join in a covalent
bond the atoms with the fewest valence
electrons will pair up with the atom that has
the most up to a total of 8.
If a pair of electrons is not shared they are
called an unshared pair.
If two atoms share two pairs of electrons it
is called a double covalent bond.
If two atoms share three pairs of electrons it
is called a triple covalent bond.

The Lewis structures that you have seen
so far use a pair of dots to represent
covalent bonds.
 Another kind of Lewis structure uses
dashes to represent covalent bonds.
– A single dash represents a single
covalent bond.
– A double dash represents a double
covalent bond
– A triple dash represents a triple
covalent bond.
 Covalent compounds (molecules) tend
to have low melting and boiling points
 When an electron pair is shared by the
direct overlap of bonding orbitals, a
sigma bond results. Single bonds are
sigma bonds.
 The overlap of parallel orbitals forms a
pi bond. Multiple bonds involve both
sigma and pi bonds.
 Bond length depends on the sizes of the
bonded atoms and the number of electron
pairs they share.
 Bond dissociation energy is the energy
needed to break a covalent bond.
– Double bonds have larger bond
dissociation energies than single. Triple
even larger
 Bond length and bond dissociation energy
are directly related.
9.2 Naming Molecules
–
Covalent compounds (Molecules) are named by adding
prefixes to the element names.
•
Covalent’ (in this context) means both elements are nonmetals.
–
The names of molecular compounds will include prefixes
that indicate the number of atoms in the molecule. (CO2
would be called Carbon Dioxide, CCl4 is called Carbon
Tetrachloride)
–
The following prefixes are used:
mono-1
tri-3
penta-5
hepta-7
nona-9
di-2
tetra-4
hexa-6
octa-8
deca-10
Note: When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in
the prefix is dropped.
– Remember to use the “ide” suffix.
(Cl is chloride not chlorine)
– Exceptions to the rule:
•
The prefix “mono” is usually not written with the
first word of the compounds name.
– CO2 is not called monocarbon dioxide, just carbon
dioxide.
•
Some compounds have common names. (H2O is
water or dihydrogen monoxide)
Naming Binary Covalent Compounds:
Examples
1
mono
2
di
3
tri
4
tetra
5
penta
6
hexa
7
heptaa
8
octa
9
nona
10
deca
* Second element
in ‘ide’ from
* Drop –a & -o
before ‘oxide’
N2S4
dinitrogen tetrasulfide
NI3
nitrogen triiodide
XeF6
xenon hexafluoride
CCl4
carbon tetrachloride
P2O5
diphosphorus pentoxide
SO3
sulfur trioxide
Naming Acids
 An acid is a molecular substance that
dissolves in water to produce hydrogen
ions (H+).
– Acids behave like an ionic compound
because when they dissolve in water,
they create cations and anions.
– The cation is always H+ and the anion
depends on the particular acid.
• Example: When the acid HCL dissolves in water it forms
H+ cations and CL- anions.
• HNO3 acid will form H+ and NO3- ions.
Naming Acids
 The name of an acid usually results from the name of the
anion(-).
– For acids, where the anion ends with the suffix
“ide”(Chloride) the name of the acid begins with the prefix
hydro.
– The acids name also includes the root name of the anion(-)
and the word acid.
– In addition you need to change the suffix of the anion from
“ide” to “ic”.
• By these rules HCL is named hydrochloric acid.
 Other acids are named without the prefix “hydro”. The acid
HNO3- derives its name from NO3- which is nitrate. HNO3 is
named Nitric Acid.
 The following are names and formulas of common acids:
Naming Acids










Anion
F-, Fluoride
CL-, Chloride
Br-, Bromine
I-, iodide
S2-, sulfide
NO3-, nitrate
CO3-2, carbonate
SO4-2, sulfate
PO4-3, phoshate
C2H3O2- , acetate
Corresponding Acid
HF, hydrofluoric Acid
HCL, hydrochloric acid
HBr, hydrobromic acid
HI, hyroiodic acid
H2S, hydrosulfuric acid
HNO3, nitric acid
H2CO3, carbonic acid
H2SO4, sulfuric acid
H3PO4, Phosphoric acid
HC2H3O2, acetic acid
9.3 Molecular Structures
 Structural formulas (Lewis Structures)
use letter symbols and bonds to show
the position of atoms.
 Drawing Lewis Structures:
Water
H
O
Each hydrogen has 1 valence
electron
and wants 1 more
The oxygen has 6 valence
electrons
and wants 2 more
They share to make each other
“happy”
Water
 Put the pieces together
 The first hydrogen is happy
 The oxygen still wants one more
HO
Water
 The second hydrogen attaches
 Every atom has full energy levels
HO
H
Multiple Bonds
 Sometimes atoms share more than one
pair of valence electrons.
 A double bond is when atoms share two
pair (4) of electrons.
 A triple bond is when atoms share three
pair (6) of electrons.
Carbon dioxide
C
O
 CO2 - Carbon is central
atom ( I have to tell you)
 Carbon has 4 valence
electrons
 Wants 4 more
 Oxygen has 6 valence
electrons
 Wants 2 more
Carbon dioxide
 Attaching 1 oxygen leaves the oxygen 1
short and the carbon 3 short
CO
Carbon dioxide

Attaching the second oxygen leaves
both oxygen 1 short and the carbon 2
short
OC O
Carbon dioxide

The only solution is to share more
O CO
Carbon dioxide

The only solution is to share more
O CO
Carbon dioxide

The only solution is to share more
O CO
Carbon dioxide

The only solution is to share more
O C O
Carbon dioxide

The only solution is to share more
O C O
Carbon dioxide

The only solution is to share more
O C O
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond

O C O
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond
8 valence
electrons

O C O
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond
8 valence
electrons

O C O
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond
8 valence
electrons

O C O
How to draw them
 To figure out if you need multiple bonds
 Add up all the valence electrons.
 Count up the total number of electrons to
make all atoms happy.
 Subtract.
 Divide by 2
 Tells you how many bonds - draw them.
 Fill in the rest of the valence electrons to
fill atoms up.
Examples
N
H
 NH3
 N - has 5 valence electrons
wants 8
 H - has 1 valence electrons
wants 2
 NH3 has 5+3(1) = 8
 NH3 wants 8+3(2) = 14
 (14-8)/2= 3 bonds
 4 atoms with 3 bonds
Examples
 Draw in the bonds
 All 8 electrons are accounted for
 Everything is full
H
H NH
Examples








HCN C is central atom
N - has 5 valence electrons wants 8
C - has 4 valence electrons wants 8
H - has 1 valence electrons wants 2
HCN has 5+4+1 = 10
HCN wants 8+8+2 = 18
(18-10)/2= 4 bonds
3 atoms with 4 bonds -will require
multiple bonds - not to H
HCN
 Put in single bonds
 Need 2 more bonds
 Must go between C and N
HC N
HCN
Put in single bonds
 Need 2 more bonds
 Must go between C and N
 Uses 8 electrons - 2 more to add

HC N
HCN
Put in single bonds
 Need 2 more bonds
 Must go between C and N
 Uses 8 electrons - 2 more to add
 Must go on N to fill octet

HC N
Another way of indicating
bonds
 Often use a line to indicate a bond
 Called a structural formula
 Each line is 2 valence electrons
H O H =H O H
Structural Examples
 C has 8 electrons
because each
line is 2 electrons
 Ditto for N
H C N
 Ditto for C here
H
C O  Ditto for O
H
Coordinate Covalent Bond
 A Coordinate Covalent Bond is when one
atom donates both electrons in a covalent
bond.
 Carbon monoxide
 CO
CO
Coordinate Covalent Bond
When one atom donates both electrons
in a covalent bond.
 Carbon monoxide
 CO

C O
Coordinate Covalent Bond
When one atom donates both electrons
in a covalent bond.
 Carbon monoxide
 CO

C O
C
O
Resonance
 Resonance is a condition when more
than one dot diagram with the same
connections is possible.
 Resonance structures differ only in the
position of the electron pairs, never the
atom positions.
9.4 Molecular Shape (VSEPR)
 Valence Shell Electron Pair Repulsion.
 Predicts three dimensional geometry of
molecules.
 The name tells you the theory.
 Valence shell - outside electrons.
 Electron Pair repulsion - electron pairs try
to get as far away as possible.
 Can determine the angles of bonds.
 And the shape of molecules
VSEPR
 Based on the number of pairs of
valence electrons both bonded and
unbonded.
 Unbonded pair are called lone pair.
 CH4 - draw the structural formula
VSEPR
H
H C H
H
 Single bonds fill
all atoms.
 There are 4 pairs
of electrons
pushing away.
 The furthest they
can get away is
109.5º.
4 atoms bonded
 Basic shape is
tetrahedral.
 A pyramid with a
triangular base.
 Same basic shape
for everything
with 4 pairs.
H
H
109.5º
C
H
H
3 bonded - 1 lone pair
 Still basic tetrahedral but you can’t see
the electron pair.
 Shape is called
trigonal pyramidal.
H N H H
H
N
H
H
<109.5º
2 bonded - 2 lone pair
 Still basic tetrahedral but you can’t see
the 2 lone pair.
 Shape is called
bent.
H O
H
O
H
H
<109.5º
3 atoms no lone pair
 The farthest you can the electron pair
apart is 120º.
 Shape is flat and called
trigonal planar.
 Will require 1 double bond
120º
H
H
H
C O
H
C
O
2 atoms no lone pair
 With three atoms the farthest they can
get apart is 180º.
 Shape called linear.
 Will require 2 double bonds or one triple
bond
180º
O C O
Molecular Orbitals
 The overlap of atomic orbitals from
separate atoms makes molecular
orbitals
 Each molecular orbital has room for two
electrons
 Two types of MO
– Sigma ( σ ) between atoms
– Pi ( π ) above and below atoms
Sigma bonding orbitals
 From s orbitals on separate atoms
+
+
s orbital s orbital
+ +
+ +
Sigma bonding
molecular orbital
Sigma bonding orbitals
 From p orbitals on separate atoms


p orbital


p orbital


Sigma bonding
molecular orbital
Pi bonding orbitals






 P orbitals on separate atoms
Pi bonding
molecular orbital
Sigma and pi bonds
 All single bonds are sigma bonds
 A double bond is one sigma and one pi
bond
 A triple bond is one sigma and two pi
bonds.
Hybridization
 The mixing of several atomic orbitals to
form the same number of hybrid orbitals.
 When an atom approaches another atom to
form a bond, the orbitals of its electrons
may be changed.
 The changing of the orbitals forms what is
called hybrid orbitals. This is a cross
between all of the orbitals involved (s,p,d,f).
Hybridization
 When an “s” and “p” orbital come together they
form a linear shaped molecule.
 When an “s” and “p2” orbital come together they
form a trigonal planer molecule.
 When an “s” and “p3” orbital come together they
form tetrahedral, pyramidal or bent molecules.
 Atoms with higher orbital numbers form longer
bonds. Multiple bonds are shorter than single
bonds because bonds with more electrons attract
the nuclei of the bonding atoms more strongly.
9.5 Electronegativity and Bond Type
Electro negativity
Difference
Bond Type
<.4
Non-Polar Covalent
0.5 – 1.9
Polar Covalent
>2.0
Ionic
Polar Bonds
 When the atoms in a bond are the
same, the electrons are shared equally.
 This is a nonpolar covalent bond.
 When two different atoms are
connected, the electrons may not be
shared equally.
 This is a polar covalent bond.
Electronegativity
 A measure of how strongly the atoms
attract electrons in a bond.
 The bigger the electronegativity
difference the more polar the bond.
0.0 - 0.4 Covalent nonpolar
0.5 - 1.0 Covalent moderately polar
1.0 -2.0 Covalent polar
>2.0 Ionic
How to show a bond is polar
 Isn’t a whole charge just a partial charge
 d+ means a partially positive
 d- means a partially negative
d+
H
d-
Cl
 The Cl pulls harder on the electrons
 The electrons spend more time near the Cl
Polar Molecules
 Polar molecules have a partially
positive end and a partially negative end
 Requires two things to be true:
– The molecule must contain polar
bonds.This can be determined from
differences in electronegativity.
– Symmetry can not cancel out the
effects of the polar bonds.
Must determine geometry first.
Polar Molecules
 Symmetrical shapes are those without
lone pair on central atom
– Tetrahedral
– Trigonal planar
– Linear
 Will be nonpolar if all the atoms are the
same
 Shapes with lone pair on central atom
are not symmetrical
 Can be polar even with the same atom
Intermolecular Forces
 They are what make solid and liquid
molecular compounds possible.
 The weakest are called van der Waal’s
forces - there are two kinds
– Dispersion forces
– Dipole Interactions
Dispersion Force
 Depends only on the number of
electrons in the molecule
 Bigger molecules more electrons
 More electrons stronger forces
• F2 is a gas
• Br2 is a liquid
• I2 is a solid
Dipole interactions
 Occur when polar molecules are
attracted to each other.
 Slightly stronger than dispersion forces.
 Opposites attract but not completely
hooked like in ionic solids.
Dipole interactions
 Occur when polar molecules are
attracted to each other.
 Slightly stronger than dispersion forces.
 Opposites attract but not completely
hooked like in ionic solids.
+
d
d
H F
+
d
d
H F
Properties of Molecular
Compounds
 Made of nonmetals
 Poor or nonconducting as solid, liquid or
aqueous solution
 Low melting point
 Two kinds of crystals
– Molecular solids held together by IMF
– Network solids- held together by bonds
– One big molecule (diamond, graphite)