Chapter 8: Covalent Bonding

Matter takes many forms in nature: In this
chapter, we are going to learn to distinguish
the type of compound that we have already
studied, the “ionic compound” (which
contains oppositely-charged particles: metal
cations and non-metal anions), from a different
type of compound – a “molecular compound”.
Additionally, we are going to focus on a type of
molecular compound known as a binary
molecular compound
II. Binary compounds:
A “binary” compound contains atoms from
two different elements.
 A. NaCl”, “CaF2”, and “Al2O3” (3 ionic
compounds) are binary ionic compounds.
“NH4Cl” is an ionic compound, but because
it contains more than 2 elements, it is not a
binary ionic compound.

B. “N2O5”, “SF6”, and H2O” (3 molecular
compounds) are binary molecular
compounds.
 “C6H12O6” is a molecular compound, but
because it contains more than 2 elements it
is not a binary molecular compound.

Comparison/Contrast between an ionic
compound and a molecular substance.
 A. Molecular substances are made of
molecules.
 1. There is no “molecule” in an ionic
compound.

B. A “molecule” contains a specific
number of atoms, connected in a specific
manner, to give a specific shape.
 If even one atom is “missing” or “different”,
the molecule would be an entirely different
substance.

1. Not so with an ionic compound: In an
ionic compound there is a specific ratio of
atoms.. In salt (NaCl) for example, there is
a ratio of 1 Na for every 1 Cl.
 If a clump of salt lost 1 Na and 1 Cl, it
would still be the same original substance:
NaCl

C. The formula of a “molecule” should
never be simplified. C2H8 is not the same
substance as CH4.
 1. The formula of an ionic compound
should always be simplified. Ba2O2 is the
same substance as BaO.
 D. A molecule will not crack apart.
 1. Ionic compounds can crack apart if
hammered….. If the cations come close too
close together and the anions come too close
together, the structure cracks apart.

E. Molecular substances may have low
melting points.
 1. All ionic compounds have a very high
melting point.
 F. Molecular substances may, at room
temperature, be found as solids, liquids, or
gases.


1. All ionic compounds are solids (at room
temperature).
G. Molecular substances contain atoms
which are held together by covalent bonds.
 1. Ionic substances are held together by
ionic bonds.

IV. Covalent bonding:
The type of bonding that occurs within a
molecular substance, in which atoms share
their valence electrons in order to become
more stable.
 A. Occurs between atoms of nonmetallic
elements.

B. Not all “molecules” or “molecular
substances” are compounds!
 In addition to the binary molecular
compounds that we will study, there are 7
nonmetallic elements found in nature (in
their elemental form) as pairs of atoms.
These are the 7 “diatomic” elements:
 N2, O2, F2, Cl2,
Br2, I2, H2.

V. An important review:
A. Metallic elements: Found to the left side
of the staircase boundary on the periodic
table.
 Non-metallic elements:
 Elements found to the right side of the
staircase boundary on the periodic table.
 2. Hydrogen is a nonmetallic element also.

VI. The octet rule:
When a molecule is formed: “Nonmetal
atoms share electrons in covalent bonds in
order to obtain a full octet of electrons.” An
octet = 8 valence electrons.
 A. Exception: A hydrogen atom will end up
with a total of two electrons by sharing with
1 other atom.

B. There are a few other notable exceptions
to the octet rule:
 1. A few molecular compounds which
contain an odd number of valence electrons
are known to exist.
 2. A few molecular compounds have either
a boron or an aluminum atom with 6
valence electrons.

2. A few molecular compounds have a
central atom with 10 or 12 valence
electrons.
 (1) One common example is “sulfur
hexafluoride”.
 In this compound, the central sulfur atom
contains 6 x 2 = 12 valence electrons. Be
sure to remember that this compound is an
example in which the central atom does not
follow the octet rule.

VII. Types of covalent bonds.

A. Single covalent bond – 1 shared pair of
valence electrons: 2 dots, or a single dash,
represent 2 electrons that are simultaneously
being attracted by, or “shared” by, the nuclei
of two neighboring atoms.
1. The formula in the center is a type of
structural formula called a “Lewis dot
structural formula”.
 The formula on the right is the molecular
formula.

H–H
H H
H2
B. Double covalent bond
 – two pairs of shared valence electrons: 4 dots
or 2 parallel dashes.

H
H
C
H
C
H
H
C
H
C
H
H
C 2H 4

C. Triple covalent bond– three pairs of
shared valence electrons: 6 dots or 3 parallel
dashes.
N

N
or
N
N
N2
Notice the two “unshared pairs” of electrons
(one pair is to the far left and one pair to the
far righ)t of the nitrogen structure. You
may never use a long dash to represent an
unshared pair of electrons.

Unshared pairs of electrons don’t bond the
atoms together….but, the repulsive forces of
unshared pairs of electrons do dramatically
influence the shape of a molecule!

D. Notice how an ion can react with a
molecule to generate a polyatomic ion. In
the example below, a hydrogen ion bonds to
a molecule of ammonia(NH3) to make the
ammonium ion (NH4)+:
H
H+ +
N H
H
[H
+
H
N
H
H
]
VIII. Drawing a Lewis Dot
Structure:
A. Certain elements are known as “central”
atoms…. They will be found in the center of
a structure. The first element given in a
formula is usually the central atom
(exception: hydrogen and the halogens).
 1. Position the central atom in the center of
your work space.

B. Hydrogen and the halogens are known
as “peripheral” atoms. They will be found
only connected to one other atom.
 Position hydrogen and halogen atoms so
that they “touch”, or “go around” only 1
other atom.

C. Add up all the valence electrons.
Position the valence electrons as dots
around the atom they belong to - the valence
electrons may never leave the original atom.
 Position the dots to form a “doorway” with
4 sides, in which the symbol of the element
appears centered in the doorway.
 Start with no more than 2 dots on each side
of the 4 sided doorway.

D. If you can’t easily achieve a Lewis dot
structure which has each atom (other than
hydrogen) surrounded by 8 dots by doing
what is described above, then you either
need a double bond (2 pairs of shared
electrons) or a triple bond (3 pairs of shared
electrons).
 For CO2, you will need two double bonds.

1. To make a double bond, move one “unshared electron” simultaneously from each
of two neighboring atoms, and place those 2
electrons in between the two neighboring
atoms.
 2. To make a triple bond, start with a
double bonded pair of atoms, and
simultaneously move one more unshared
electron from each of the two atoms.
Reposition those two electrons in between
the atoms.


Important points regarding nonmetal atoms
and their bonding charcteristics:
Atoms of the
Following
Nonmetallic
Elements:
Have This
Number of
Electrons
(dots) when
in a Stable
Structure:
Type of Bonds
Permitted by these
atoms:
Things to Remember
about atoms of these
elements; or to remember
about a specified
molecule.
Hydrogen
2
Single covalent
Diatomic element;
H2 is a “linear” diatomic
molecule
Boron,
Aluminum
8 or 6
Single covalent
When only 6 electrons
surround a boron or an
aluminum atom, the
molecule’s shape will be
“trigonal planar” (a flat
pancake).
Atoms of the
Following
Nonmetallic
Elements:
Sulfur
The Halogen
family: F, Cl,
Br, I, At
Have This
Number of
Electrons
(dots) when
in a Stable
Structure:
Type of Bonds
Permitted by these
atoms:
Things to Remember
about atoms of these
elements; or to remember
about a specified
molecule.
8 EXCEPT
with
“SF6” when
there are 12
Single, double,
and/or triple
covalent
In sulfur hexafluoride
sulfur does NOT follow
the octet rule. This is
one “exception” to the
octet rule.
Single covalent
All are diatomic
elements; and,
F2, Cl2, Br2, I2, At2 are
all linear molecules,
with only single bonds.
8
Atoms of the
Following
Nonmetallic
Elements:
Have This
Type of Bonds
Number of
Permitted by these
Electrons
atoms:
(dots) when in
a Stable
Structure:
Things to Remember about
atoms of these elements; or
to remember about a
specified molecule.
The Noble Gas
family:
8….Except Do NOT form
for helium compounds easily
(2). Noble (no bonds).
gas atoms
don’t form
compounds.
Always found as single
atoms in the gaseous
state.
Nitrogen and
oxygen
8
Single, double,
and/or triple
covalent
Diatomic elements; linear
molecules. N2 has one
triple bond, while O2 has
1 double bond.
All other nonmetal
atoms
8
Single, double,
and/or triple
covalent
IX. Lewis Dot structural formulas for
polyatomic ions:
 A. Covalent bonds occur within a
polyatomic ion (not between polyatomic
ions).
 B. When drawing polyatomic ions, place
the first element in the center of the
structure, and place the second element
around the first element (placing 1 atom of
the second element along each different side
of the first element).



C. When the charge of a polyatomic ion is +,
you need to subtract the indicated number of
electrons from the total of the valence electrons
in the molecule. So, for +1 ions: take away 1
electron from the molecular ion’s number of
valence electrons.
D. When the charge of a polyatomic ion is –,
you need to add the indicated number of
electrons to the molecular ion’s number of
valence electrons.
1. If the charge is 1-, then add 1 more
electron to the molecule’s total number of
valence electrons.
 2. If the charge is 2-, then add 2 more
electrons; if the charge is 3-, then add 3
more electrons.
 E. Last, for a polyatomic ion: Draw a large
bracket around the ion; and, place its
charge at upper right.

# of N valence electrons: 5
# of H valence electrons: 4 x 1 = 4
Charge of ion = +1, therefore less 1
Therefore, total =
Ammonium ion (NH4)+
[
H
H N H
H
5
4
-1
8
+
]

Steps for Dot Structures:

Step 1: total # valence electrons.

Step 2. Position central atoms:
 carbon
atoms form a straight line;
 assume only single bonding.

Step 3. Position other atoms; remember “special”
molecules.

A. Peripheral atoms:
 Hydrogen and the halogens connect to only 1 other atom,
 use only 1 single bond.

B. binary polyatomic ions: first element is central, second
element is peripheral. assume all single bonds.

C. hydrocarbons” – molecular formula gives list of atoms
(from left to right) connecting to each central atom (usually
carbon)


CH3CH2OH means “first carbon touches 3 H atoms, second
carbon touches 2 H atoms, then there is an O touching an H.
Assume all single bonds.
D. Memorize: CO2 C in the middle; use two double bonds
•
Step 4: Make each atom stable.
Work from left to right:
Assume all single bonds.
Position unshared pairs to provide octets.
Exception: Hydrogen atoms = only 2 dots.
Step 5: Count the dots you’ve used.
Make sure the # you used = the # you were
supposed to. Erase extras.

Step 6: Make corrections 
If your structure “needs” 2 extra dots, it really
needs a double bond….
 Erase
2 unshared dots, and share them (as part of
a double bond).

If your structure “needs” 4 extra dots, it really
needs a triple bond.....
 Erase
4 unshared dots, and share them (as part of
a triple bond).

Every atom should now be stable.
X. VSEPR Theory –
Valence Shell Electron Pair Repulsion theory.
[Remember: Like charges repel!]
 A. A theory to predict the 3-dimensional
geometry, ie. the“shape” of a molecule

1. The theory is based on “electrostatic
repulsion”: Molecules will adjust their
shape to keep the negatively-charged pairs
of valence electrons as far apart as possible
from each other.
 B. When NOT to use VSEPR theory: When
there are only 2 atoms in a molecule. These
molecule’s shapes are called linear – it
doesn’t matter if there are single bonds,
double bonds, triple bonds, or unshared
electron pairs.

C. Using VSEPR theory:
1. Draw the Lewis dot structure for the
molecule.
 2. Identify its central atom.
 3. Identify the sets of valence electrons as one
of two possibilities:
 A. Those connecting two atoms.
 B. Those that do not connect two atoms.
These are called “unshared pairs”.

4. The unshared pairs found on a central atom
strongly repel each other; and molecules that
would otherwise be linear, will be forced into a
bent (or angular) shape.
 5. Unshared pairs also cause a molecule
that would be shaped like a flat triangle
(trigonal planar), to be forced into a not flat
( trigonal pyramidal) shape.

6.Count the number of connections separately
from the number of unshared pairs.
 1 single bond counts as 1 connection.
 1 double bond counts as 1 connection.
 1 triple bond counts as 1 connection.
 Each unshared set of 2 dots counts as 1 unshared pair.

D Predicting Shapes Using VSEPR
Table

Read horizontally across the table.
Connections To
the Central Atom
Unshared Pairs of
Molecular Shape
Electrons Around Central
Around Central Atom
Atom
2
0
Linear
3
0
Trigonal Planar
4
0
Tetrahedral, 109.5o
2
1 or 2
Bent
3
1
Pyramidal
Shapes:
,
Linear diatomic
E. Shapes:
Trigonal planar
Trigonal Planar
Bent
Linear diatomic
Linear triatomic
Linear triatomic
Pyramidal
Tetrahedral
Trigonal pyramidal
tetrahedral
bent
“Molecular Polarity”
– A term that is used to distinguish two types
of molecules…. Based on the presence or
absence of a separation of charge.
 Some molecules show characteristics
indicating that they have oppositely-charged
ends (a positive end and a negative end).
This is called a separation of the charges (or
“separation of charge”).


Other molecules show characteristics
indicating that their structure doesn’t have
a separation of charge, or their structure
hides the presence of their oppositelycharged ends.
How to determine a molecule’s
polarity.
The first part of determining a molecule’s
polarity is to calculate each individual bond’s
polarity.
 Be careful with the vocabulary being used –
 An individual bond’s polarity is called the
“bond polarity”
 The polarity of the entire molecule is called
the “molecular polarity”

To calculate a bond polarity, first identify
the “electronegativity value” of each of the 2
atoms in the bond you are working on.
 The electronegativity value is number (from
0 to 4) which informs us of an atom’s ability
to attract electrons when in a compound.




The electronegativity value is given on your
periodic table, side 2, within each element’s
square…..upper right corner of the square, in
black print.
The closer an element’s electronegativity is to
“4, the better that an atom of that element will
attract electrons when that atom is found in a
compound.
The closer an element’s electronegativity is to
“0”, the less likely it is for that atom to be able
to attract electrons when in a compound.

After identifying the two electronegativity
values, subtract the two electronegativity
values. Take the absolute value of the
answer (make the answer positive).

After subtracting and taking the absolute value of the two
electronegativity values, then think about your answer, and
determine whether your answer indicates that there is, or that there
is not, a situation in which one of the two atoms in the bond
“overpowers” the other atom in terms of electron attracting ability.
If an atom is able to overpower the other atom, it will “hog” the
electrons, as opposed to sharing the electrons equally with the
other atom. You will see (on the next page) that I’ve placed a
“δ–”sign next to an oxygen atom, and a “δ+” sign next to 2
hydrogen atoms, in a sketch of a water molecule. By doing this,
I am indicating that the oxygen atom is hogging the negativelycharged valence electrons belonging to the 2 hydrogen atoms;
and, the 2 hydrogen atoms are both overpowered and “partially
lose” their own valence electrons to the oxygen atom.

δ- indicates the “partial negative” atom; and,
δ+ indicates the “partial positive” atom.





Finally, you are now able to conclude that a bond is
either “nonpolar covalent” or “polar covalent”.
A bond is to be called “non-polar covalent” when the two
atoms share the electrons more-or-less equally.
Labeling a bond as nonpolar covalent means that the
difference in electronegativity values fell between 0 and 0.4
Labeling a bond as nonpolar covalent means that the
electrons are distributed practically equally along the bond.
Labeling a bond as nonpolar covalent means that there is
no separation of charge in that particular bond.



A bond is to be called “polar covalent” when the
electrons are NOT distributed equally along the
bond, rather, the electrons are found much or most of
the time toward the atom that has the higher
electronegsativity value.
This will occur when difference in electronegativity
values is greater than 0.4, but less than 2.
We place a δ – next to an atom to identify it as the
“partial negative” atom; and, we place a δ + next to
an atom to identify it as the “partial positive” atom in
a polar covalent bond.



A bond is designated as ionic when one atom has
stripped the other of some of its valence electrons,
yielding a cation and an anion. Then, the cation and
anion are attracted, and more cations and anions
are attracted and an ionic compound forms.
This should occur when the difference in
electronegativity values is greater than 2.
We studied this type of compound previously. Ionic
compounds have very different characteristics than
compounds in which polar covalent and nonpolar
covalent bonds exist.


The second part of determining the “molecular
polarity” (a molecule’s polarity) is visualizing the
effect of the individual bond polarities in
conjunction with the shape of the molecule.
Some common occurrences are listed below:
If the molecule’s shape is either tetrahedral,
linear triatomic or trigonal planar and if all
the bonds in the structure are identical, then
the overall molecule’s polarity (the molecular
polarity) is nonpolar


This is because, if every bond is nonpolar,
there isn’t a substantial separation of charge to
begin with.
And, if every bond is polar (as long as each
bond has the same 2 atoms), the polarity of
each bond will be cancelled due to the
symmetry of the polar bonds around a central
atom. When this symmetry exists, the δ –
charge(s) cancel the δ + charge(s) in each bond.


If the molecule’s shape is either tetrahedral ,
trigonal planar, or linear, and the structure contains
non-identical bonds in which 1 is polar, then the
overall (molecular) polarity is polar (because 1 polar
bond won’t have another bond to be cancelled with).
If the molecules’ shapes are pyramidal or bent, and
one or more of the bonds is polar, the overall
(molecular) polarity is “polar”. This is because in
molecules having these shapes, the bond polarity is
not situated symmetrically relative to the molecule’s
center, so the polarity doesn’t cancel.
Polar bonds NonPolar molecule
Polar bonds  Polar molecule.
 Polar bonds combine to cancel out;
Polar bonds do NOT cancel out;
 Producing Non-polar molecules:
Producing Polar Molecules:

“Like dissolves like”:

An expression stating that ionic compounds
and molecules whose molecular polarity is
polar will be able to dissolve only in solvents
containing polar molecules; and, conversely
molecules whose molecular polarity is
nonpolar will be able to dissolve only in
solvents containing nonpolar molecules.
Some characteristics of a water
molecule:
A H2O molecule has 2 identical polar
bonds, a bent structure, and its molecular
polarity is “polar”.
 The partial negative (or slightly negative)
region of a water molecule is the area closest
to the oxygen atom.
 The partial positive (or slightly positive)
region of a water molecule lies within the
area closest to the two hydrogens atoms.

Because of having the above characteristics,
water behaves in an unusual manner:
 When placed between a positively charged
metal plate and a negatively charged metal
plate, the water molecules all line up, with
their positively-charged region attracted
toward the negative plate and their
negatively- charged region attracted toward
the positive plate.

Inter- molecular Bonds (or forces of
attraction)
are DIFFERENT from the covalent bonds
you have been studying so far!
 The covalent bonds you have been studying
so far (single, double, triple bonds) are
bonds within a molecule; these would be
called______ intra-molecular bonds.
 Inter-molecular bonds (forces) are the
attractions between 2 molecules. They hold
separate molecules together.

They are weaker than polar and non-polar
covalent bonds (single/double/triple).
 They are weaker than the ionic bonds which
connect cations to anions.
 Breaking an inter-molecular bond is a
physical change; whereas, breaking an
intra-molecular bond is a chemical change.


When you boil water (to generate water
vapor), or melt ice (to generate liquid
water), you are breaking inter-molecular
bonds. When you treat a water molecule
with electricity, you destroy the water
molecule and generate from it oxygen gas
and hydrogen gas….this is the breaking of
intra-molecular bonds….this is a chemical
change.
Types of intermolecular bonds:
Hydrogen bond: Is the strongest
intermolecular force; and, it is perhaps the
most important intermolecular bond, as it is
necessary for life as we know it.
 Hydrogen bonds take effect when you have
a combination of a few certain atoms in a
polar molecule.

The polar molecule must contain a
hydrogen atom, and that hydrogen atom
must be connected to one of the highly
electronegative atoms listed below:
 Fluorine, oxygen, OR nitrogen.
 Ex: H2O, HF, NH3

Hydrogen bonds determine the properties
of water and biological molecules (such as
proteins).
 Cause water to predominate as a liquid
(rather than as a gas) on earth.
 Cause ice to expand upon freezing (rather
than contract as the kinetic molecular
theory would predict).
 Holds the DNA double helix structure
together.


Time-permitting, we will also learn about
one of the weaker types of intermolecular
forces of attraction by doing a lab in which
we use evaporation rates to identify between
hydrogen bonding and a weaker force,
present between nonpolar molecules.