Chapter 8: Covalent Bonding
The Covalent Bond
• Atoms bond to gain stability and to have
access to 8 valence electrons (octet rule).
• A covalent bond arises by the sharing of
valence electrons between nonmetal atoms.
• A molecule is formed when two atoms bond
covalently.
Diatomic Molecules
• Seven elements exist as a pair of atoms, not just a
single atom. This way they are more stable. They
are called diatomic molecules.
• You need to memorize these elements:
–
–
–
–
–
–
–
Nitrogen (N2)
Oxygen (O2)
Fluorine (F2)
Chlorine (Cl2)
Bromine (Br2)
Hydrogen (H2)
Iodine (I2)
Single Covalent Bonds
• A single covalent bond occurs when only one
pair of electrons is shared between atoms.
• The shared electron pair is called a bonding
pair of electrons.
Multiple Covalent Bonds
• Double bonds are formed from the sharing of
two pairs (4 total) electrons between two
atoms.
• Triple bonds are formed from the sharing of
three pairs (6 total) electrons between two
atoms.
Rules For Naming
Covalent (or Binary Molecular)
Compounds
1. The first element is named first, using the
entire element name
2. The second element is named, using the root
of the element name plus the suffix –ide.
3. Prefixes are used to indicate the number of
atoms of each element present in the
compound.
– **Exception – THE FIRST ELEMENT NEVER USES
THE PREFIX “MONO”!
Common Prefixes
Number of
Atoms
1
Prefix
Prefix
Mono-
Number of
atoms
6
2
Di-
7
Hepta-
3
Tri-
8
Octa-
4
Tetra-
9
Nona-
5
Penta-
10
Deca-
Hexa-
Examples of Naming
Covalent Compounds
• P2O5
– There are 2 phosphorus atoms (shown by the
subscript)
– There are 5 oxygen atoms (also shown by the
subscript)
– Name the first element with its prefix
• Diphosphorus
– Name the second element with its prefix and the
suffix –ide
• Pentoxide (not pentaoxide, because of the two vowels)
– Put them together  diphosphorus pentoxide
Some Examples
•
•
•
•
•
•
•
•
CCl4
Carbon tetrachloride (not monocarbon…)
N2O
Dinitrogen monoxide
NF3
Nitrogen trifluoride
CO
Carbon monoxide
Writing formulas from names
• Use the prefixes to identify the subscripts
• Example: carbon dioxide has 1 carbon and 2
oxygens
So, it would be written as
CO2
Some examples
•
•
•
•
•
•
•
•
Triphosphorus Pentachloride
P3Cl5
Dinitrogen Trioxide
N2O3
Carbon dioxide
CO2
Dihydrogen monosulfide
H2S
Properties of Covalent Compounds
• Covalent compounds have weak forces
between molecules so the bonds break more
easily. (They are called intermolecular forces)
– Low melting points
– Low boiling points
• However, Covalent network solids are formed
in some substances which lead to very strong
solids (quartz and diamond are examples)
Lewis Structures
• Recall that electron dot structures are used to
show valence electrons.
• Unpaired electrons in the electron dot
structures are able to create covalent bonds
with other atoms.
• In a Lewis structure, they can represent the
arrangement of electrons in a molecule.
Some electron dot structures
and single covalent bonds
• Group 17 – have seven valence electrons, so one
is unpaired – these elements can make one single
covalent bond
• Group 16 – have six valence electrons, so two are
unpaired – these elements can make 2 single
covalent bonds
• Group 15 – have five valence electrons, so three
are unpaired – these elements can make 3 single
covalent bonds.
• Group 14 – have four valence electrons, all of
which are unpaired – these elements can make 4
single covalent bonds
MOLECULAR SHAPES
• The VSEPR Theory
• Valence Shell Electron Pair Repulsion Theory
• Electron pairs orient themselves in order to
minimize repulsive forces.
Molecular Shapes
• Your book shows the orbital shapes
• Go to Glencoe.com for an interactive version
• http://glencoe.mcgrawhill.com/sites/007874637x/student_view0/ch
apter8/concepts_in_motion.html
Electronegativity and Polar Bonds
• Electronegativity is an atoms “affinity” for
electrons, or its ability to draw electrons to it
• If there is a difference in electronegativity
between atoms, you will have an unequal
sharing of electrons, which is called a polar
covalent bond
•
mg
Electronegativity Differences
and Bond Type
Electronegativity
Difference
> 1.7
0.4-1.7
<0.4
0
Bond Type
Mostly ionic
Polar covalent
Mostly covalent
Nonpolar covalent