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Chapter 8
Covalent bonding
Valence Electrons
 Elements with similar chemical behavior have the
same number of valence electrons.
 For the representative elements (1A, 2A, 3A, 4A, 5A, 6A, 7A,
8A) the group number corresponds to the number of
valence electron in each group (with the exception of He)
 When examining electron configurations, the
electrons that are present in the highest principle
energy level represent the valence electrons of those
atoms.
Br: [Ar] 4s2 3d10 4p5
Bromine has 7 valence electrons
Valence Electrons and Electron Dot
Structures
 Valence electrons are the electrons that participate in
chemical bonds
 Electron dot structures consist of the atom symbol and its
valence electrons represented as dots.
Br: [Ar]4s2 3d10 4p5
Covalent Bonding
 Covalent bonds occur between two or more non-metals
 Unlike ionic bonds where electrons are transferred from
one atom to another, electrons are shared between atoms
in a covalent bond.
 Atoms joined together by covalent bonds are called
molecules
 A compound composed of molecules is called a molecular
compound
Molecular and Structural Formulas
 A molecular formula indicates the types and numbers
of each atom in a molecule
 The structural formula indicate the arrangement of
the atoms in the molecule
H2O
Covalent Bonds and the Octet Rule
 Atoms share electrons in a covalent bond so that each
atom has enough electrons to satisfy the octet rule
Varieties of Covalent Bonds
Single bonds (sigma bonds)
 One pair of electrons is shared between two atoms
 Overlap of orbitals occurs between the two nuclei
Lone pair
Varieties of Covalent Bonds
Double Bonds (1sigma bond, 1 pi bond)
 Atoms share two pairs of electrons
 Overlap of orbitals in the pi bond occur
off to the side of adjoining nuclei
Triple Bonds (1sigma bond, 2 pi bonds)
 Atoms share three pairs of electrons
Coordinate Covalent Bonds
 A covalent bond in which one atom contributes both
bonding electrons.
Carbon
Resonance Structures
A condition when more than one valid Lewis
structure can be written for a molecule or ion.
Exceptions to the Octet Rule
 Too few electrons surrounding the central atom (ex: BH3)
Boron will not have a full octet, only 6 electrons. It can only achieve a full
octet when another atom shares an entire pair of electrons with it
(Coordinate covalent bonding)
 Too many electrons surrounding the central atom (ex: PCl5)
 An odd number of electrons
How to Draw a Lewis Structure for Molecules

Predict the location of atoms
1.
If there are more than two atoms, place the least electronegative atom in the
center and surround it by the remaining atoms.
2.
Hydrogen is always terminal (outside) because it can only make one bond

Determine the total number of electrons if each atom had a full set of valence
electrons (2 for H, 8 for all others)

Add up the number of valence electron that you have to work with

Subtract total valence electrons from total electrons and divide by two. This is
the number of bonding pairs that are needed to put together the molecule.

Connect the atoms with the number of bonds that you calculated above

Add lone pairs where needed so that each atom has a full octet (except for
hydrogen which can only have two electrons)
Molecule
HCN
Total Electrons
Valence Electrons Bonding Pairs
Polyatomic Ions
Polyatomic ions are a cluster of non-metals that carry a charge.
To draw the structure of a polyatomic ion, follow the procedure
for drawing ordinary molecules but add or subtract the
number of electrons gained or lost to the total number of
valence electrons in your structure as indicated by the charge
on the ion.
Molecule
IO3-
Total Electrons
Valence Electrons
Bonding Pairs
Molecular Shape (VSEPR)
Valence Shell Electron Pair Repulsion – minimizes the repulsion of shared
and unshared pairs of electrons around the central atom.
 The shape of a molecule determines many of its physical and
chemical properties.
 The VSEPR is based on the arrangement of bonding and lone
electrons around a central atom to minimize repulsion.
 The repulsion of electrons creates a specific bond angle between a
central atom and two terminal atoms.
 Lone pairs of electrons occupy more space than bonding pairs of
electrons
Molecular Geometry
Electronegativity and Polarity
Recall: Electronegativity is the ability of an atom to attract an electron.
Chemical bonding is like “Tug-o-War”
Bond Type
Electronegativity
Difference
Non-polar Covalent
Polar Covalent
Ionic
0-0.4
0.5-2.0
>2.0
Molecular Polarity
Molecules are either polar or non-polar
Both polar and non-polar molecules may contain polar bonds. What determines
whether a molecule is polar or non-polar is the symmetry of the molecule
PolarBonds
Present
Symmetr Polar/
y
NonPolar
Examples
No
No
Non-Polar
NO2
No
Yes
Non-polar
SiH4
Yes
No
Polar
NH3
Yes
Yes
Non-polar
CO2
VSEPR shapes that can demonstrate symmetry are:
Linear
Trigonal Planar
Tetrahedral
Naming Binary Covalent Compounds
• If there is more than one of the 1st atom, precede the atom name by the appropriate
prefix
(di, tri, tetra, penta, hexa, hepta, octa, nona, deca)
Example: C6O2
hexacarbon dioxide
•If there is only one of the first atom, do not precede the atom name by mono.
CO2 = monocarbon dioxide
CO2 = carbon dioxide
•Precede the second atom name by the appropriate prefix, including mono if there is
only one of that atom. Drop the last syllable (or 2) and add –ide to the element name
C2O Dicarbon monoxide
2nd Element
C
N
O
F
P
Name
Carbide
Nitride
Oxide
Fluoride
Phosphide
2nd Element
S
Cl
Se
Br
I
Name
Sulfide
Chloride
Selenide
Bromide
Iodide